Atomic Structure and the Periodic Table

Historical Models of the Atom

Have you ever wondered how scientists figured out what matter is made of? Everything around us, including air, water, rocks, plants, and people, is made of tiny particles called atoms. Today, scientists know a lot about atoms, but that understanding took many years to build.

Scientists did not discover the atom all at once. Instead, different scientists made models to explain what atoms might look like and how they behave. As new experiments were done, the old models changed. This is an important idea in science: scientific models improve when new evidence is found.

In this lesson, you will learn about four important historical models of the atom:

• Dalton's solid sphere model
• Thomson's plum pudding model
• Rutherford's nuclear model
• Bohr's model

By the end, you should be able to explain how each model is different and why scientists changed their ideas over time.

1. Dalton's Solid Sphere Model

In the early 1800s, John Dalton helped develop one of the first modern atomic theories. He said that all matter is made of tiny particles called atoms.

Dalton thought atoms were like tiny, solid balls. In his model, atoms could not be broken into smaller parts. This is why his idea is often called the solid sphere model.

Dalton also believed:

• All atoms of the same element are alike.
• Atoms of different elements are different.
• Atoms join together to form compounds.
• In chemical changes, atoms are rearranged, not created or destroyed.

Dalton's model was an important starting point because it helped scientists explain why substances combine in certain ways. Even though we now know atoms are made of smaller parts, Dalton's idea that matter is made of atoms was a big step forward.

Picture Dalton's model like this: a tiny, hard, solid ball with no smaller pieces inside.

2. Thomson's Plum Pudding Model

Later, in the late 1800s, J. J. Thomson discovered that atoms contain even smaller particles called electrons. Electrons have a negative charge. This discovery showed that atoms were not solid balls after all.

Since atoms are usually neutral, Thomson reasoned that there must also be positive charge in the atom to balance the negative electrons. He suggested a new model called the plum pudding model.

In this model:

• The atom was a sphere of positive matter.
• Negative electrons were scattered throughout it.

The name comes from a dessert. Imagine a pudding with small pieces of fruit mixed inside. In Thomson's model, the electrons were like the fruit, and the positive part of the atom was like the pudding.

This model was important because it introduced the idea that atoms have smaller parts inside them.

3. Rutherford's Nuclear Model

In the early 1900s, Ernest Rutherford tested Thomson's model with an experiment. He shot tiny charged particles at a thin piece of gold foil. Most of the particles passed straight through, but a few bounced back.

This result was surprising. If Thomson's model were correct, the particles should have mostly gone through with only small changes in direction.

Rutherford concluded that:

• Most of the atom is empty space.
• Almost all of the atom's mass is packed into a tiny center called the nucleus.
• The nucleus has a positive charge.
• Electrons move around outside the nucleus.

This new idea became the nuclear model of the atom.

Picture Rutherford's model like this: a tiny, dense center in the middle, with electrons outside it and lots of empty space in between.

This was a major change from Thomson's model. The positive charge was not spread out through the whole atom. Instead, it was concentrated in the nucleus.

4. Bohr's Model

Niels Bohr improved Rutherford's model. He studied how electrons behave and suggested that electrons do not move just anywhere around the nucleus. Instead, they move in certain paths or energy levels around the nucleus.

This became known as the Bohr model.

In Bohr's model:

• The nucleus is at the center of the atom.
• Electrons move around the nucleus in fixed energy levels.
• Electrons can move from one energy level to another.

A simple way to picture this is to think of the energy levels like rings around the nucleus.

Bohr's model helped explain why atoms give off light in certain ways. It also gave students and scientists a clearer picture of how electrons are arranged.

Why the Models Changed

Each model of the atom changed because scientists found new evidence. Science is not just guessing. Scientists make observations, do experiments, and then improve their explanations.

Here is the pattern scientists followed:

1. Make a model to explain what is known.
2. Test the model with experiments.
3. Find new evidence.
4. Change the model if needed.

This means newer models are usually more accurate than older ones.

Comparing the Models

• Dalton: Atom is a solid, unbreakable sphere.
• Thomson: Atom has electrons mixed into a positive sphere.
• Rutherford: Atom has a small, positive nucleus and mostly empty space.
• Bohr: Electrons move in fixed energy levels around the nucleus.

Simple Timeline

1. Dalton — atoms are solid spheres
2. Thomson — atoms contain electrons
3. Rutherford — atoms have a nucleus
4. Bohr — electrons are in energy levels

Worked Example 1: Identifying Dalton's Model

Question: A student says, "An atom is a tiny, solid ball with no smaller parts inside." Which scientist's model matches this idea?

Step 1: Look for the key clue. The atom is described as a solid ball.

Step 2: Think about which scientist described atoms this way.

Answer: This matches Dalton's solid sphere model.

Why: Dalton believed atoms were tiny, solid particles that could not be divided.

Worked Example 2: Identifying Thomson's Change

Question: What important new idea did Thomson add to atomic theory?

Step 1: Think about what Dalton's model was missing.

Step 2: Remember that Thomson discovered electrons.

Answer: Thomson showed that atoms contain smaller parts inside them, especially negatively charged electrons.

Why: This proved atoms were not just solid, featureless spheres.

Worked Example 3: Understanding Rutherford's Evidence

Question: In Rutherford's gold foil experiment, most particles passed through the foil, but a few bounced back. What did this tell scientists?

Step 1: If most particles passed through, then much of the atom must not block them.

Step 2: If a few bounced back, there must be a small, dense part that can strongly push them away.

Answer: The atom is mostly empty space, with a small, dense, positively charged nucleus in the center.

Why: A spread-out positive atom would not cause some particles to bounce straight back.

Worked Example 4: Comparing Rutherford and Bohr

Question: Both Rutherford and Bohr said atoms have a nucleus. What did Bohr add to Rutherford's model?

Step 1: Start with what they had in common: a nucleus at the center.

Step 2: Ask what Bohr said about electrons.

Answer: Bohr said electrons move in fixed energy levels around the nucleus.

Why: Rutherford showed the nucleus exists, but Bohr gave a clearer idea of how electrons are arranged around it.

Tips for Remembering the Models

• Dalton = solid ball
• Thomson = electrons in pudding
• Rutherford = tiny nucleus, empty space
• Bohr = rings or energy levels

Why This Matters

Learning about historical models of the atom helps us understand how science works. Scientific ideas are built step by step. When better evidence is found, scientists improve their explanations.

These atomic models also help us understand the modern study of matter, the periodic table, and how elements behave.

Brief Summary

The atomic model changed over time as scientists learned more. Dalton said atoms were solid spheres. Thomson discovered electrons and proposed the plum pudding model. Rutherford showed that atoms have a tiny, positive nucleus and are mostly empty space. Bohr then explained that electrons move in fixed energy levels around the nucleus.

When you study these models, focus on what each scientist added. That makes it easier to see how atomic theory developed from a simple solid sphere to a more detailed picture of the atom.

Subatomic Particles are the tiny parts that make up an atom. Everything around you is made of atoms, and every atom is built from even smaller pieces. The three main subatomic particles are protons, neutrons, and electrons.

Learning about these particles helps us understand what atoms are like, why different elements are different, and how atoms behave.

What is an atom?

An atom is the basic building block of matter. Matter is anything that has mass and takes up space. A chair, a drop of water, and even the air are all made of atoms.

Atoms are extremely small, but they still have parts inside them. These parts are the subatomic particles.

The three subatomic particles

• Protons have a positive charge.
• Neutrons have no charge. They are neutral.
• Electrons have a negative charge.

Charge is a way to describe a kind of electric property. Positive and negative charges can balance each other.

Where are these particles found?

Protons and neutrons are found in the nucleus, which is the center of the atom. The nucleus is small and dense.

Electrons are found outside the nucleus in the space around it. You can think of them as moving around the nucleus in an electron cloud.

• Nucleus: contains protons and neutrons
• Outside the nucleus: contains electrons

Comparing mass

Protons and neutrons each have about the same mass. Electrons have much less mass than either protons or neutrons.

For 6th grade science, it helps to remember their relative mass like this:

• Proton: about 1
• Neutron: about 1
• Electron: about 0 compared to a proton or neutron

This does not mean an electron has no mass. It means its mass is so small that we often treat it as almost zero when comparing it to the other two particles.

A quick comparison chart

• Proton: positive charge, mass about 1, located in the nucleus
• Neutron: no charge, mass about 1, located in the nucleus
• Electron: negative charge, mass about 0, located outside the nucleus

Why protons matter

The number of protons in an atom tells you what element it is. For example, every hydrogen atom has 1 proton. Every carbon atom has 6 protons. Every oxygen atom has 8 protons.

If the number of protons changes, the atom becomes a different element.

Why neutrons matter

Neutrons add mass to the atom. They also help make the nucleus more stable.

Atoms of the same element can have different numbers of neutrons, but they are still the same element because the number of protons stays the same.

Why electrons matter

Electrons are involved in how atoms interact with each other. They are important in chemical bonding, which is how atoms join together.

In a neutral atom, the number of electrons is equal to the number of protons. This is because the positive charge of the protons balances the negative charge of the electrons.

For a neutral atom:

$$\text{number of protons} = \text{number of electrons}$$

How charges add up

Each proton has a charge of \(+1\). Each electron has a charge of \(-1\). Each neutron has a charge of \(0\).

To find the total charge of an atom, you can combine the charges from protons and electrons.

For example, if an atom has 3 protons and 3 electrons:

$$(+3) + (-3) = 0$$

That atom is neutral.

If an atom has 5 protons and 4 electrons:

$$(+5) + (-4) = +1$$

That atom has an overall positive charge.

Worked Example 1: Identifying each particle

Question: Which subatomic particle has a negative charge and is found outside the nucleus?

Step 1: Think about the charges.

• Proton = positive
• Neutron = neutral
• Electron = negative

Step 2: Think about the location.

• Protons and neutrons are in the nucleus.
• Electrons are outside the nucleus.

Answer: The particle is the electron.

Worked Example 2: Comparing mass and location

Question: Which two subatomic particles have about the same mass, and where are they found?

Step 1: Recall relative mass.

• Proton = about 1
• Neutron = about 1
• Electron = about 0

Step 2: Find the particles with matching masses.

Protons and neutrons both have about the same mass.

Step 3: Recall their location.

Both are found in the nucleus.

Answer: Protons and neutrons have about the same mass, and they are found in the nucleus.

Worked Example 3: Finding whether an atom is neutral

Question: An atom has 8 protons, 8 neutrons, and 8 electrons. Is it neutral?

Step 1: Look at protons and electrons.

Protons are positive, and electrons are negative.

Step 2: Compare their numbers.

The atom has 8 protons and 8 electrons.

Step 3: Decide whether the charges balance.

Since the number of protons equals the number of electrons, the charges cancel out.

$$(+8) + (-8) = 0$$

Answer: Yes, the atom is neutral.

Worked Example 4: Describing an atom from its particles

Question: A particle in an atom has no charge and is in the nucleus. What is it, and what is its relative mass?

Step 1: Find the particle with no charge.

The particle with no charge is the neutron.

Step 2: Recall its location and mass.

A neutron is found in the nucleus and has a relative mass of about 1.

Answer: It is a neutron, and its relative mass is about 1.

Easy way to remember the three particles

• Proton: positive, center, mass 1
• Neutron: neutral, center, mass 1
• Electron: negative, outside center, mass 0

You can also remember that p in proton goes with positive.

Common mistakes to avoid

• Do not confuse neutron with electron. A neutron has no charge and is in the nucleus. An electron has a negative charge and is outside the nucleus.
• Do not say electrons are in the nucleus. They are outside it.
• Do not forget that protons and neutrons have much more mass than electrons.
• Do not forget that the number of protons identifies the element.

Brief summary

Atoms are made of three main subatomic particles: protons, neutrons, and electrons. Protons are positive, neutrons are neutral, and electrons are negative.

Protons and neutrons are found in the nucleus and each have a relative mass of about 1. Electrons are found outside the nucleus and have much less mass, about 0 compared with protons and neutrons.

In a neutral atom, the number of protons equals the number of electrons. Understanding these particles helps explain the structure of atoms and the differences between elements.

Atomic Number and Mass Number

Everything around us is made of tiny particles called atoms. Atoms are the building blocks of matter. Even though atoms are very small, they are made of even smaller parts.

The three main parts of an atom are protons, neutrons, and electrons. For this lesson, we will focus mostly on protons and neutrons, because they help us find the atomic number and the mass number.

Where are these parts found?

• Protons are in the center of the atom, called the nucleus.
• Neutrons are also in the nucleus.
• Electrons move around outside the nucleus.

What is the atomic number?

The atomic number tells how many protons are in an atom. This number is very important because it tells us which element the atom is.

For example, every atom with 1 proton is hydrogen. Every atom with 6 protons is carbon. Every atom with 8 protons is oxygen. If the number of protons changes, the element changes.

So, the atomic number is like an element's ID number.

Atomic number = number of protons

We can write that as:

$$\text{Atomic Number} = \text{Number of Protons}$$

What is the mass number?

The mass number tells the total number of protons and neutrons in the nucleus.

To find the mass number, add the number of protons and the number of neutrons.

Mass number = protons + neutrons

We can write that as:

$$\text{Mass Number} = \text{Protons} + \text{Neutrons}$$

Electrons are not included in the mass number.

Why not count electrons?

At this level, you can think of electrons as having so little mass compared with protons and neutrons that we do not include them when finding the mass number.

Important difference

• Atomic number = only protons
• Mass number = protons + neutrons

Many students mix these up, so it helps to remember:

• Atomic number identifies the atom.
• Mass number measures how many heavy particles are in the nucleus.

Using the atomic number to identify an element

The periodic table lists all the elements. Each element has its own atomic number. If you know the atomic number, you can find the element.

• Atomic number 1 = Hydrogen
• Atomic number 2 = Helium
• Atomic number 6 = Carbon
• Atomic number 7 = Nitrogen
• Atomic number 8 = Oxygen
• Atomic number 11 = Sodium

This means if an atom has 8 protons, it must be oxygen. It does not matter how many neutrons it has. The number of protons decides the element.

How to find the number of neutrons

If you know the mass number and the atomic number, you can find the number of neutrons by subtracting.

$$\text{Neutrons} = \text{Mass Number} - \text{Atomic Number}$$

This works because the atomic number is the number of protons.

So if:

• mass number = protons + neutrons
• atomic number = protons

then subtracting gives the neutrons.

Step-by-step method

1. Find the atomic number. That tells you the number of protons.
2. If needed, use the periodic table to identify the element.
3. Find the mass number.
4. Add protons and neutrons to check mass number, or subtract to find neutrons.

Worked Example 1: Finding the atomic number

An atom has 7 protons. What is its atomic number?

Step 1: Remember that atomic number equals number of protons.

$$\text{Atomic Number} = \text{Number of Protons}$$

Step 2: Substitute 7 for the number of protons.

$$\text{Atomic Number} = 7$$

Answer: The atomic number is 7.

Worked Example 2: Finding the mass number

An atom has 6 protons and 8 neutrons. What is its mass number?

Step 1: Use the formula.

$$\text{Mass Number} = \text{Protons} + \text{Neutrons}$$

Step 2: Substitute the numbers.

$$\text{Mass Number} = 6 + 8$$

$$\text{Mass Number} = 14$$

Answer: The mass number is 14.

Worked Example 3: Finding neutrons

An atom has an atomic number of 11 and a mass number of 23. How many neutrons does it have?

Step 1: Atomic number 11 means the atom has 11 protons.

Step 2: Use the neutrons formula.

$$\text{Neutrons} = \text{Mass Number} - \text{Atomic Number}$$

Step 3: Substitute the numbers.

$$\text{Neutrons} = 23 - 11$$

$$\text{Neutrons} = 12$$

Answer: The atom has 12 neutrons.

Worked Example 4: Identifying the element and mass number

An atom has 8 protons and 10 neutrons.

• What is its atomic number?
• What is its mass number?
• What element is it?

Step 1: Find the atomic number.

Atomic number = number of protons

$$\text{Atomic Number} = 8$$

Step 2: Find the mass number.

$$\text{Mass Number} = \text{Protons} + \text{Neutrons}$$

$$\text{Mass Number} = 8 + 10 = 18$$

Step 3: Identify the element.

An atom with atomic number 8 is oxygen.

Answer:

• Atomic number = 8
• Mass number = 18
• Element = oxygen

Common mistakes to avoid

• Do not confuse atomic number with mass number.
• Do not add electrons when finding mass number.
• Do not forget that the element is decided by protons.
• When finding neutrons, subtract: mass number - atomic number.

Quick practice ideas

• If an atom has 3 protons, its atomic number is 3.
• If an atom has 4 protons and 5 neutrons, its mass number is 9.
• If an atom has atomic number 12 and mass number 24, it has 12 neutrons.

Brief Summary

The atomic number is the number of protons in an atom. It tells which element the atom is.

The mass number is the total number of protons and neutrons. You can use these formulas:

$$\text{Atomic Number} = \text{Protons}$$

$$\text{Mass Number} = \text{Protons} + \text{Neutrons}$$

$$\text{Neutrons} = \text{Mass Number} - \text{Atomic Number}$$

If you remember what each number counts, you can identify elements and solve atom problems with confidence.

Organization of the Periodic Table

The periodic table is a chart that organizes all known elements. It helps scientists and students see patterns in the way elements are arranged and how they behave.

An element is a pure substance made of only one kind of atom. Each box on the periodic table shows one element, such as hydrogen, oxygen, or gold.

Learning how the periodic table is organized makes it much easier to understand science. You can use the table to find an element, compare it to others, and notice which elements have similar properties.

How the periodic table is arranged

The periodic table is arranged in order of atomic number. The atomic number tells how many protons are in the nucleus of an atom.

For example, hydrogen has atomic number 1, so it has 1 proton. Helium has atomic number 2, so it has 2 protons. As you move across the table, the atomic number increases by 1 each time.

You can think of the periodic table as being organized in two main directions:

• Rows, called periods
• Columns, called groups

Periods: the rows of the table

A period is a horizontal row across the periodic table. Elements in the same period have the same number of energy levels for their electrons.

In 6th Grade science, it is helpful to think of energy levels as layers or rings around the nucleus where electrons can be found.

Here is the main idea:

• Elements in Period 1 have 1 energy level.
• Elements in Period 2 have 2 energy levels.
• Elements in Period 3 have 3 energy levels.

This means the period number matches the number of energy levels.

For example:

• Hydrogen and helium are in Period 1, so they each have 1 energy level.
• Lithium, carbon, oxygen, and neon are in Period 2, so they each have 2 energy levels.
• Sodium, magnesium, and chlorine are in Period 3, so they each have 3 energy levels.

Groups: the columns of the table

A group is a vertical column on the periodic table. Elements in the same group often have similar chemical properties. This means they often act in similar ways during chemical reactions.

Why do they act alike? One important reason is that they have the same number of electrons in their outer energy level. These outer electrons help determine how an element reacts.

For 6th Grade, the most important idea is this: same group = similar properties.

For example, elements in the same group may:

• react in similar ways
• form similar kinds of compounds
• share other physical and chemical traits

Examples of groups

Some groups have special names. You do not need to memorize all of them, but seeing a few helps you understand the patterns.

• Group 1: includes hydrogen, lithium, sodium, and potassium. Most of these elements are very reactive.
• Group 17: includes fluorine and chlorine. These elements are also reactive.
• Group 18: includes helium, neon, and argon. These elements are very unreactive compared with many others.

Notice that elements in each column share similar behaviors. That is one of the biggest reasons the periodic table is so useful.

What information is in an element box?

Each element has its own box on the periodic table. A box often shows:

• Atomic number
• Chemical symbol
• Element name
• Atomic mass (on many tables)

For example, a box for oxygen may show:

• Atomic number: 8
• Symbol: O
• Name: Oxygen

The atomic number tells the number of protons. In a neutral atom, the number of electrons is the same as the number of protons, so oxygen also has 8 electrons.

Reading the table from left to right

As you move from left to right across a period, the atomic number increases one by one. So if one element has atomic number 11, the next one has atomic number 12.

This makes the periodic table different from a random chart. It is arranged in a careful order based on the structure of atoms.

Metals, nonmetals, and metalloids

The periodic table also helps us group elements by general type.

• Metals are mostly on the left side and center of the table.
• Nonmetals are mostly on the right side.
• Metalloids are found along a stair-step line between metals and nonmetals and have properties of both.

You do not need to memorize every element type right now, but it is helpful to know that the table shows these broad patterns too.

Why groups matter

If two elements are in the same group, they often behave in similar ways. This helps scientists predict what an unfamiliar element might do.

For example, if sodium is in the same group as lithium, and lithium is reactive, then sodium is likely to be reactive too. The periodic table gives clues before we even do an experiment.

Why periods matter

If an element is in Period 2, it has 2 energy levels. If it is in Period 4, it has 4 energy levels. This helps us understand the structure of the atom.

So groups tell us about similar properties, and periods tell us about energy levels.

A simple way to remember it

• Group = Go up and down (vertical column)
• Period = across, like a sentence period moves you along a line (horizontal row)

Another memory trick is:

• Groups share properties
• Periods show energy levels

Worked Example 1: Finding a period

Question: Neon is in Period 2. How many energy levels does neon have?

Step 1: Look at the period number.

Neon is in Period 2.

Step 2: Match the period number to the number of energy levels.

Period 2 means 2 energy levels.

Answer: Neon has 2 energy levels.

Worked Example 2: Finding a group pattern

Question: Sodium and potassium are in the same group. What can you predict about their properties?

Step 1: Remember what a group means.

Elements in the same group have similar chemical properties.

Step 2: Apply the rule.

If sodium and potassium are in the same group, they will likely react in similar ways.

Answer: Sodium and potassium probably have similar chemical properties.

Worked Example 3: Using both group and period

Question: Chlorine is in Group 17 and Period 3. What does this tell you?

Step 1: Use the group number idea.

Group 17 means chlorine has properties similar to other elements in Group 17.

Step 2: Use the period number idea.

Period 3 means chlorine has 3 energy levels.

Answer: Chlorine shares similar properties with other elements in Group 17 and has 3 energy levels because it is in Period 3.

Worked Example 4: Reading atomic number order

Question: An element has atomic number 12. The next element to the right has what atomic number?

Step 1: Remember the pattern across a period.

Atomic number increases by 1 as you move left to right.

Step 2: Add 1.

$$12 + 1 = 13$$

Answer: The next element to the right has atomic number 13.

Common mistakes to avoid

• Mixing up groups and periods: Groups go up and down. Periods go across.
• Forgetting what periods show: The period number tells the number of energy levels.
• Forgetting what groups show: Elements in the same group have similar properties.
• Thinking the table is random: The elements are arranged by increasing atomic number.

Quick check for understanding

1. If an element is in Period 4, how many energy levels does it have?
2. If two elements are in the same group, what is likely true about them?
3. Do groups go across or up and down?
4. What happens to atomic number as you move left to right?

Answers:

1. 4 energy levels
2. They likely have similar chemical properties.
3. Up and down
4. It increases by 1 each element.

Summary

The periodic table is a chart of elements arranged by increasing atomic number. Its rows are called periods, and the period number tells the number of energy levels in the atom.

Its columns are called groups, and elements in the same group usually have similar chemical properties. When you understand groups and periods, the periodic table becomes much easier to read and use.

Metals, Nonmetals, and Metalloids

Everything around us is made of elements. Scientists organize elements on the periodic table. One helpful way to understand elements is to sort them into three big groups: metals, nonmetals, and metalloids.

These groups are based on how elements behave and what properties they have. We can compare them by looking at things like conductivity (how well something carries heat or electricity), malleability (whether it can be hammered into shapes), and luster (whether it looks shiny).

Learning these groups helps us predict how an element might act. For example, if an element is a metal, it will probably be shiny and conduct electricity well. If it is a nonmetal, it will probably not conduct electricity well. If it is a metalloid, it may act like both.

1. What Are Metals?

Metals are the largest group of elements on the periodic table. Most elements are metals.

Metals usually have these properties:

• Shiny appearance, or luster
• Good conductors of heat and electricity
• Malleable, which means they can be hammered into thin sheets
• Ductile, which means they can be pulled into wires
• Usually solid at room temperature

Examples of metals include:

• Iron (Fe)
• Copper (Cu)
• Aluminum (Al)
• Gold (Au)
• Silver (Ag)

Metals are useful in everyday life. Copper is used in electrical wires because it carries electricity well. Aluminum is used in cans and foil because it is lightweight and easy to shape. Iron is used in buildings and tools because it is strong.

Many metals also tend to lose electrons when they join with other elements. This means they often donate electrons. At your grade level, it is enough to remember that metals often give away electrons more easily than nonmetals do.

2. What Are Nonmetals?

Nonmetals are a smaller group of elements. They have properties that are very different from metals.

Nonmetals usually have these properties:

• Often dull, not shiny
• Poor conductors of heat and electricity
• If solid, they are often brittle, which means they break easily
• Many are gases at room temperature
• They are not malleable or ductile

Examples of nonmetals include:

• Oxygen (O)
• Carbon (C)
• Nitrogen (N)
• Sulfur (S)
• Chlorine (Cl)

Nonmetals are also important in everyday life. Oxygen is needed for breathing. Carbon is found in living things. Chlorine is used to help keep pools clean.

Nonmetals often gain electrons or share electrons when they combine with other elements. This is one reason they behave differently from metals.

3. What Are Metalloids?

Metalloids are elements with properties of both metals and nonmetals. They are like a middle group.

Metalloids usually have these properties:

• They may look shiny like metals
• They are often brittle like nonmetals
• They may conduct electricity somewhat, but not as well as metals

Examples of metalloids include:

• Silicon (Si)
• Boron (B)
• Germanium (Ge)

Silicon is a very important metalloid. It is used in computers, phones, and calculators because it can control the flow of electricity. Materials like this are called semiconductors. For 6th Grade, you can think of a semiconductor as a material that conducts electricity better than a nonmetal, but not as well as a metal.

4. Where Are These Groups on the Periodic Table?

The periodic table helps us see where these groups are located.

• Metals are mostly on the left side and in the middle of the table.
• Nonmetals are mostly on the right side of the table.
• Metalloids are found along a zigzag stair-step line between metals and nonmetals.

A simple way to remember this is:

• Left and center = mostly metals
• Right side = mostly nonmetals
• Along the staircase = metalloids

Hydrogen is a special case. It is on the left side of the periodic table, but it is a nonmetal.

5. Comparing the Three Groups

Here is a simple comparison:

• Metals: shiny, bendable, good conductors, often donate electrons
• Nonmetals: dull, brittle if solid, poor conductors, often gain or share electrons
• Metalloids: have mixed properties, can conduct some electricity

Another way to compare them is by conductivity:

• Metals conduct electricity well
• Metalloids conduct electricity a little or under certain conditions
• Nonmetals conduct electricity poorly

You can think of this as an order from greatest to least conductivity:

$$\text{metals} > \text{metalloids} > \text{nonmetals}$$

6. Why These Groups Matter

Classifying elements helps scientists and students predict how materials will be used.

For example:

• If you need a wire, a metal is a good choice because it conducts electricity and can be pulled into thin strands.
• If you need a material for a computer chip, a metalloid may be a good choice because it can control electricity.
• If you need a gas for breathing, you would use a nonmetal like oxygen.

These categories also help explain how elements combine with each other. Metals often combine with nonmetals because metals tend to give electrons and nonmetals tend to gain or share electrons.

Worked Example 1: Classifying by Properties

Question: An element is shiny, can be hammered into a sheet, and carries electricity well. Is it a metal, nonmetal, or metalloid?

Step 1: Look at the clues. The element is shiny, malleable, and a good conductor.

Step 2: Compare with the three groups. These are all common properties of metals.

Answer: The element is a metal.

Worked Example 2: Using the Periodic Table

Question: An element is found on the right side of the periodic table. Is it most likely a metal, nonmetal, or metalloid?

Step 1: Remember the layout of the periodic table.

• Left and center = metals
• Right side = nonmetals
• Zigzag line = metalloids

Step 2: Since the element is on the right side, it is most likely a nonmetal.

Answer: It is most likely a nonmetal.

Worked Example 3: Mixed Properties

Question: An element is brittle like a nonmetal, but it can conduct electricity better than most nonmetals. What group does it belong to?

Step 1: Notice that it has a mix of properties.

Step 2: A material with both metal-like and nonmetal-like properties is usually a metalloid.

Answer: The element is a metalloid.

Worked Example 4: Choosing a Material

Question: A student wants to make a simple electrical wire for a science project. Should the student choose a metal, nonmetal, or metalloid?

Step 1: Wires need to carry electricity well.

Step 2: Wires also need to be shaped into long, thin strands.

Step 3: Metals are good conductors and are ductile, so they can be made into wires.

Answer: The student should choose a metal.

Common Mistakes to Avoid

• Do not assume all solids are metals. Some nonmetals are solid too, such as carbon and sulfur.
• Do not assume all shiny elements are metals. Some metalloids can also be shiny.
• Do not forget that hydrogen is a nonmetal even though it is placed on the left side of the table.
• Do not confuse brittle with malleable. Brittle materials break easily. Malleable materials can be hammered into shape.

Quick Check

1. Which group is usually the best conductor of electricity?
2. Which group has properties of both metals and nonmetals?
3. Which group is often dull and brittle if solid?
4. Where on the periodic table are metalloids usually found?

Answers:

1. Metals
2. Metalloids
3. Nonmetals
4. Along the zigzag stair-step line

Summary

Elements can be grouped as metals, nonmetals, or metalloids. Metals are usually shiny, bendable, and good conductors. Nonmetals are usually dull, poor conductors, and brittle if solid. Metalloids have a mix of properties and are often found along the stair-step line of the periodic table.

If you remember the main properties and where each group is located on the periodic table, you can usually classify an element correctly. This makes the periodic table easier to understand and helps explain why different materials are used for different jobs.

Covalent Bonding is a type of chemical bonding where atoms share electrons.

This usually happens between nonmetal atoms. By sharing electrons, the atoms can fill their outer energy level and become more stable.

You can think of covalent bonding like two people sharing supplies so both have what they need. Instead of one atom giving electrons away, the atoms share them.

Why do atoms bond?

Atoms bond because they are more stable when their outer energy level is full. The outer energy level is the part of the atom where bonding happens.

For many common atoms, being stable means having 8 electrons in the outer energy level. This is sometimes called the octet rule. Hydrogen is a little different and is stable with 2 electrons.

What makes covalent bonds different?

• Covalent bond: electrons are shared
• Usually forms between: nonmetals
• Forms: molecules

A molecule is a group of atoms held together by covalent bonds.

Some common molecules are:

• Water: \(H_2O\)
• Oxygen gas: \(O_2\)
• Carbon dioxide: \(CO_2\)
• Methane: \(CH_4\)

Valence electrons

The electrons in the outer energy level are called valence electrons. These are the electrons involved in bonding.

To understand covalent bonding, we look at how many valence electrons an atom already has and how many more it needs to be stable.

• Hydrogen has 1 valence electron and needs 1 more.
• Oxygen has 6 valence electrons and needs 2 more.
• Nitrogen has 5 valence electrons and needs 3 more.
• Carbon has 4 valence electrons and needs 4 more.
• Chlorine has 7 valence electrons and needs 1 more.

How sharing works

When two atoms share a pair of electrons, they form a single covalent bond. A shared pair means 2 electrons are being shared.

We can show one shared pair like this:

$$H:H$$

This means the two hydrogen atoms share 2 electrons.

Sometimes atoms need to share more than one pair of electrons.

• Single bond: 1 shared pair of electrons
• Double bond: 2 shared pairs of electrons
• Triple bond: 3 shared pairs of electrons

These can be shown with lines:

• Single bond: \(H-H\)
• Double bond: \(O=O\)
• Triple bond: \(N\equiv N\)

Lewis dot models

A simple way to model covalent bonding is with a Lewis dot model. In this model, dots show valence electrons around an element symbol.

For example:

• Hydrogen has 1 dot
• Oxygen has 6 dots
• Chlorine has 7 dots

When atoms bond, shared electrons are shown between the atoms. This helps us see how each atom reaches a full outer energy level.

Worked Example 1: Hydrogen molecule, \(H_2\)

Each hydrogen atom has 1 valence electron. Each hydrogen needs 1 more electron to have a full outer energy level of 2.

So, two hydrogen atoms share 1 pair of electrons.

Model:

$$H-H$$

Now each hydrogen counts the shared pair, so each has 2 electrons in its outer level. That makes both hydrogen atoms stable.

Worked Example 2: Oxygen molecule, \(O_2\)

Each oxygen atom has 6 valence electrons. Each oxygen needs 2 more electrons to reach 8.

That means the two oxygen atoms must share 2 pairs of electrons.

Model:

$$O=O$$

This is a double bond. Each oxygen now counts 8 electrons in its outer energy level.

Worked Example 3: Water, \(H_2O\)

Water has 2 hydrogen atoms and 1 oxygen atom.

Oxygen has 6 valence electrons and needs 2 more. Each hydrogen has 1 valence electron and needs 1 more.

Oxygen shares one pair with one hydrogen and another pair with the second hydrogen. That makes two single bonds.

Model:

$$H-O-H$$

Now each hydrogen has 2 electrons, and oxygen has 8 electrons in its outer energy level. Water is a molecule made by covalent bonding.

Worked Example 4: Carbon dioxide, \(CO_2\)

Carbon has 4 valence electrons, so it needs 4 more to reach 8. Each oxygen has 6 valence electrons, so each oxygen needs 2 more.

One carbon atom can share with two oxygen atoms. Carbon shares 2 pairs with one oxygen and 2 pairs with the other oxygen.

Model:

$$O=C=O$$

This means carbon forms a double bond with each oxygen. Now carbon and both oxygen atoms have full outer energy levels.

How to tell if covalent bonding will happen

1. Look at the elements involved.
2. If they are nonmetals, they will often form covalent bonds.
3. Check how many valence electrons each atom has.
4. See how many electrons each atom needs to fill its outer energy level.
5. Figure out how the atoms can share electrons to become stable.

Covalent bonding and formulas

A chemical formula tells how many atoms are in a molecule.

• \(H_2\) means 2 hydrogen atoms
• \(O_2\) means 2 oxygen atoms
• \(H_2O\) means 2 hydrogen atoms and 1 oxygen atom
• \(CO_2\) means 1 carbon atom and 2 oxygen atoms

The small number written lower than the symbol is called a subscript. It tells the number of atoms of that element.

Important ideas to remember

• Covalent bonds form when atoms share electrons.
• Covalent bonding usually happens between nonmetals.
• Covalent bonds form molecules.
• Valence electrons are the electrons used in bonding.
• Atoms share electrons to become more stable.
• A single bond shares 2 electrons, a double bond shares 4, and a triple bond shares 6.

Common mistakes

• Mistake: Thinking atoms in a covalent bond give away electrons.
  Fix: In covalent bonding, atoms share electrons.
• Mistake: Thinking covalent bonds usually happen between metals and nonmetals.
  Fix: Covalent bonds usually form between nonmetals.
• Mistake: Forgetting that hydrogen only needs 2 electrons, not 8.
  Fix: Hydrogen is stable with 2 electrons in its outer energy level.

Brief Summary

Covalent bonding is the sharing of electron pairs between nonmetal atoms. This sharing helps atoms fill their outer energy levels and become stable. Covalent bonds make molecules such as \(H_2\), \(O_2\), \(H_2O\), and \(CO_2\). By counting valence electrons, we can model how atoms share electrons and predict the bonds they form.

Chemical Formulas are a short way to show which elements are in a substance and how many atoms of each element are present.

When scientists write a chemical formula, they use element symbols from the periodic table. For example, hydrogen is written as H and oxygen is written as O.

A chemical formula can tell you the ratio of atoms in a compound. A compound is a substance made of two or more different elements joined together.

For example, water has the formula H₂O. This means each water particle has 2 hydrogen atoms and 1 oxygen atom.

In this lesson, you will learn how to read subscripts and coefficients so you can figure out exactly how many atoms are shown in a chemical formula.

1. Element Symbols in Formulas

Each element has a one-letter or two-letter symbol. These symbols are always written in a specific way:

• The first letter is capitalized.
• If there is a second letter, it is lowercase.

Examples:

• H = hydrogen
• O = oxygen
• Na = sodium
• Cl = chlorine

It is important to read symbols carefully. For example, CO and Co are not the same. CO means carbon and oxygen, while Co is the symbol for cobalt.

2. What Subscripts Mean

A subscript is a small number written a little lower than the symbol. It tells how many atoms of that element are in the formula.

In the formula CO₂, the 2 is a subscript. It means there are 2 oxygen atoms.

If there is no subscript after an element, that means there is 1 atom of that element.

Examples:

• H₂O = 2 hydrogen atoms and 1 oxygen atom
• CO₂ = 1 carbon atom and 2 oxygen atoms
• NaCl = 1 sodium atom and 1 chlorine atom

You can think of the formula as a recipe. The subscript tells how many of each kind of atom you need.

3. What Coefficients Mean

A coefficient is a regular-sized number placed in front of a chemical formula. It tells how many whole molecules or units of that substance there are.

For example, in 2H₂O, the coefficient is 2. This means there are 2 water molecules.

Since each water molecule has 2 hydrogen atoms and 1 oxygen atom, 2 water molecules have:

•  \(2 \times 2 = 4\) hydrogen atoms
•  \(2 \times 1 = 2\) oxygen atoms

So 2H₂O means 4 hydrogen atoms and 2 oxygen atoms total.

A coefficient multiplies every atom in the formula after it.

4. Subscripts and Coefficients Together

To read a formula correctly:

1. Find the coefficient, if there is one.
2. Read each element symbol.
3. Use the subscript to find how many atoms of each element are in one molecule.
4. Multiply by the coefficient if needed.

This can be written as:

$$\text{total atoms} = \text{coefficient} \times \text{subscript}$$

If there is no coefficient, use 1. If there is no subscript, use 1.

For example, in 3CO₂:

• Coefficient = 3
• Carbon has no subscript, so carbon = 1
• Oxygen has subscript 2, so oxygen = 2

Now multiply:

• Carbon: \(3 \times 1 = 3\)
• Oxygen: \(3 \times 2 = 6\)

So 3CO₂ means 3 carbon atoms and 6 oxygen atoms total.

5. Reading Common Chemical Formulas

Here are some common formulas and what they mean:

• O₂ = 2 oxygen atoms
• N₂ = 2 nitrogen atoms
• CH₄ = 1 carbon atom and 4 hydrogen atoms
• NH₃ = 1 nitrogen atom and 3 hydrogen atoms
• CaCl₂ = 1 calcium atom and 2 chlorine atoms

Notice that formulas can have more than two kinds of atoms, but the same rules still work.

Worked Example 1: Reading a Simple Formula

What does H₂O mean?

Step 1: Identify the elements: H is hydrogen, O is oxygen.

Step 2: Read the subscripts.

• Hydrogen has a subscript of 2, so there are 2 hydrogen atoms.
• Oxygen has no subscript, so there is 1 oxygen atom.

Answer: H₂O means 2 hydrogen atoms and 1 oxygen atom.

Worked Example 2: Formula with Two Subscripts

What does CO₂ mean?

Step 1: Identify the elements: C is carbon, O is oxygen.

Step 2: Read the subscripts.

• Carbon has no subscript, so there is 1 carbon atom.
• Oxygen has a subscript of 2, so there are 2 oxygen atoms.

Answer: CO₂ means 1 carbon atom and 2 oxygen atoms.

Worked Example 3: Using a Coefficient

How many atoms are in 2H₂O?

Step 1: The coefficient is 2, so there are 2 water molecules.

Step 2: One water molecule has:

• 2 hydrogen atoms
• 1 oxygen atom

Step 3: Multiply by the coefficient.

• Hydrogen: \(2 \times 2 = 4\)
• Oxygen: \(2 \times 1 = 2\)

Answer: 2H₂O has 4 hydrogen atoms and 2 oxygen atoms.

Worked Example 4: A More Challenging Formula

How many atoms of each element are in 3CH₄?

Step 1: The coefficient is 3, so there are 3 molecules of CH₄.

Step 2: In one molecule of CH₄:

• Carbon = 1 atom
• Hydrogen = 4 atoms

Step 3: Multiply each by 3.

• Carbon: \(3 \times 1 = 3\)
• Hydrogen: \(3 \times 4 = 12\)

Answer: 3CH₄ has 3 carbon atoms and 12 hydrogen atoms.

6. Common Mistakes to Avoid

• Do not ignore missing subscripts. If there is no subscript, the number of atoms is 1.
• Do not let the coefficient change only one element. The coefficient multiplies all atoms in the formula.
• Read symbols carefully. Capital and lowercase letters matter.
• Do not confuse formulas with words. The formula shows exact numbers of atoms.

7. Quick Check Tips

When you see a chemical formula, ask yourself:

1. What are the element symbols?
2. What does each subscript say?
3. Is there a coefficient in front?
4. What is the total number of atoms of each element?

If you answer these questions in order, you can understand most 6th grade chemical formulas.

Summary

Chemical formulas use element symbols and numbers to show the kinds and numbers of atoms in a substance.

Subscripts tell how many atoms of one element are in a molecule. If there is no subscript, it means 1.

Coefficients tell how many whole molecules there are, and they multiply all the atoms in the formula.

By reading symbols carefully and using subscripts and coefficients, you can find the exact ratio and total number of atoms in a compound.