5 - Basic Principles of Matter and Energy

Question 1

Define:

atom

Answer 1

The smallest unit of an element that retains its chemical properties.

Atoms consist of a nucleus (protons and neutrons) and electrons that orbit the nucleus.

Question 2

Explain:

What are the three subatomic particles of an atom?

Answer 2

1. Protons
2. Neutrons
3. Electrons

Protons: Positively charged particles located in the nucleus.

Neutrons: Neutral particles with no charge, also found in the nucleus.

Electrons: Negatively charged particles that orbit the nucleus.

Question 3

True or False:

Electrons are located inside the nucleus of an atom.

Answer 3

False

Electrons orbit the nucleus in distinct energy levels, while protons and neutrons are in the nucleus.

Question 4

Identify:

What forms when two or more atoms bond?

Answer 4

molecule

Molecules can consist of identical atoms (O₂) or different atoms (H₂O).

Question 5

Identify:

What is the charge of a neutron?

Answer 5

It has no charge; it is electrically neutral.

Neutrons are found in the nucleus of an atom, alongside protons, and contribute to the atom’s mass but not its electrical charge.

Question 6

Identify:

List the characteristics of ions.

Answer 6

• Have a net electrical charge.
• Formed by the loss or gain of electrons.
• Can be positive (cation) or negative (anion).

Ions play a critical role in chemical bonding, conductivity, and biological functions.

Question 7

Describe:

How are elements different from compounds?

Answer 7

• Elements are pure substances of one atom type.
• Compounds are two or more elements chemically bonded.

For example, Oxygen (O₂) is an element, and water (H₂O) is a compound.

Question 8

Identify:

Term for a positively charged ion.

Answer 8

cation

Cations form when atoms lose electrons, commonly seen in metals like sodium (Na⁺).

Question 9

Explain:

Why do atoms form chemical bonds?

Answer 9

To have a stable and complete outer electron shell.

Chemical bonds include ionic, covalent, and metallic types, each with unique properties.

Question 10

True or False:

All compounds are molecules, but not all molecules are compounds.

Answer 10

True

Every compound is a molecule because it consists of two or more atoms bonded together. However, not all molecules are compounds, as some are made of only one type of element (e.g., O₂ or N₂).

Question 11

Identify:

What forms when substances combine without changing their properties?

Answer 11

mixture

Mixtures can be separated by physical means, unlike compounds.

Question 12

Explain:

What are the 2 main types of mixtures?

Answer 12

1. Homogeneous
2. Heterogeneous

Homogeneous: Uniform composition (e.g., saltwater).

Heterogeneous: Non-uniform composition (e.g., sand in water).

Homogeneous mixtures have indistinguishable parts, while heterogeneous mixtures do not.

Question 13

Identify:

Examples of homogeneous mixtures.

Answer 13

• Saltwater
• Air
• Vinegar

Homogeneous mixtures have a uniform composition, meaning their individual components cannot be easily distinguished from each other.

Question 14

Explain:

Why is air considered a homogeneous mixture?

Answer 14

Because its gases are evenly mixed and cannot be visually separated.

Air contains nitrogen, oxygen, and other gases in uniform proportions.

Question 15

Identify:

Examples of heterogeneous mixtures.

Answer 15

• Salad
• Oil and water
• Granite

Heterogeneous mixtures have components that are not uniformly distributed and can be easily identified and separated.

Question 16

Describe:

How is a solution different from a suspension?

Answer 16

• A solution is evenly mixed.
• A suspension has particles that settle.

Saltwater is a solution; muddy water is a suspension.

Question 17

Identify:

What forms when a substance exceeds its solubility in a solution?

Answer 17

precipitate

Precipitates are solid particles formed in a liquid solution, often used to identify specific ions or compounds.

Question 18

Define:

solubility

Answer 18

The ability of a substance to dissolve in a solvent.

Solubility is influenced by factors such as temperature, pressure, and the nature of the solute and solvent (e.g., sugar dissolves more readily in warm water).

Question 19

Identify:

What is the solute in a sugar-water solution?

Answer 19

sugar

The solute dissolves in the solvent; in this case, water is the solvent.

Question 20

Identify:

Substance that dissolves the solute in a solution.

Answer 20

solvent

Water, the universal solvent, dissolves a wide variety of substances.

Question 21

Explain:

What are common methods for separating mixtures?

Answer 21

• Filtration
• Distillation
• Evaporation
• Chromatography

Filtration: Separates solids from liquids.

Distillation: Separates substances by boiling points.

Evaporation: Removes a liquid, leaving solid solutes behind.

Chromatography: Separates components based on movement through a medium.

Question 22

Describe:

How is solubility affected by temperature?

Answer 22

It increases with temperature for solids but decreases for gases.

This principle explains why warm water dissolves more sugar, but carbonated drinks lose fizz when warm.

Question 23

Define:

solid

Answer 23

A state of matter with a definite shape and volume.

Solids have tightly packed particles that vibrate in fixed positions.

Question 24

Identify:

Which state of matter has a definite volume but no fixed shape?

Answer 24

liquid

Liquids flow to take the shape of their container but maintain a constant volume.

Question 25

# Explain: Do gases have a **definite volume**?

Answer 25

No, they expand to fill the container they are in, meaning their volume is determined by the container and not a fixed property of the gas itself.

Question 26

# Describe: How do particles in a **plasma** differ from those in a **gas**?

Answer 26

* Plasma has **ionized** particles. * Gas has **neutral** particles. ## Footnote Plasma conducts electricity due to its charged particles, unlike gases.

Question 27

# Identify: Which state of matter has the **highest** energy level?

Answer 27

plasma ## Footnote **Plasma** contains charged particles and is found in stars and lightning.

Question 28

# Explain: Why do solids **maintain** their shape?

Answer 28

Because their particles are **tightly packed** and **fixed**. ## Footnote The strong intermolecular forces in solids prevent changes in shape under normal conditions.

Question 29

# Explain: Why can't liquids be **compressed** easily?

Answer 29

* **Close molecules**: Liquid molecules are already packed tightly. * **Incompressible molecules**: The molecules themselves resist compression. * **Fixed volume**: Liquids maintain a specific volume. ## Footnote Liquids are nearly **incompressible** due to the close arrangement of their particles.

Question 30

# Identify: **Characteristics** of a liquid.

Answer 30

* Takes the shape of its container. * Has definite volume. * Particles can move and slide past each other. ## Footnote **Liquids** are fluid and adapt to their container's shape while maintaining volume.

Question 31

# Define: sublimation

Answer 31

When a solid changes **straight** into a gas without melting. ## Footnote An example is **dry ice** (solid CO₂), which sublimates at room temperature.

Question 32

# Describe: How is evaporation **different** from boiling?

Answer 32

**Evaporation** happens at the surface, while **boiling** occurs at a specific temperature. ## Footnote Both are phase changes where liquids become gases, but boiling requires reaching the boiling point.

Question 33

# Identify: The phase change where a gas **turns** into a liquid.

Answer 33

condensation ## Footnote **Condensation** occurs when gas particles lose energy, as seen with dew formation.

Question 34

# Define: chemical property

Answer 34

A **property** that appears only in a **chemical reaction**. ## Footnote **Examples:** flammability, reactivity, and the ability to oxidize. These properties involve changes in a substance's chemical structure.

Question 35

# Identify: What are common **indicators** of a chemical change?

Answer 35

* Formation of a precipitate. * Change in color. * Release of gas (bubbling). * Temperature change. ## Footnote These changes suggest the formation of a **new substance**, as seen in reactions like rusting or vinegar reacting with baking soda.

Question 36

# Identify: What property can be **observed without changing** a substance’s composition?

Answer 36

physical property ## Footnote Common physical properties include *color*, *density*, *melting point*, and *boiling point*.

Question 37

# Identify: What is the **pH range** of acidic substances?

Answer 37

0 to 6.9 ## Footnote Acidity *increases* as pH decreases. Substances like lemon juice (pH ~2) are strongly acidic, while milk (pH ~6) is weakly acidic.

Question 38

# Identify: What is a substance with a pH of **9** considered?

Answer 38

Basic (alkaline) ## Footnote **Basic** substances, such as baking soda and ammonia, neutralize acids and often have slippery textures.

Question 39

# Define: oxidation

Answer 39

Chemical process in which a **substance loses electrons**, often involving the addition of oxygen or the removal of hydrogen. ## Footnote While it often involves oxygen, it can also occur in reactions without it, such as in the rusting of metals.

Question 40

# Explain: What are the main **types** of chemical bonds?

Answer 40

* Covalent Bonds * Ionic Bonds * Metallic Bonds ## Footnote **Covalent Bonds**: Electrons are shared between atoms. **Ionic Bonds**: Electrons are transferred from one atom to another, forming charged ions. **Metallic Bonds**: A sea of delocalized electrons surrounds positive metal ions.

Question 41

# Identify: What type of bond is **formed** by the sharing of electrons between atoms?

Answer 41

covalent bond ## Footnote *Covalent bonds* occur in molecules like water (H₂O), where atoms share electrons to achieve stability.

Question 42

# Identify: 3 examples of **chemical changes**.

Answer 42

1. Burning wood 1. Rusting iron 1. Digesting food ## Footnote In chemical changes, bonds are broken or formed, resulting in new substances with different properties.

Question 43

# Identify: 3 examples of **physical changes**.

Answer 43

1. Melting ice 1. Boiling water 1. Breaking glass ## Footnote Physical changes only affect the form of a substance, not its chemical structure.

Question 44

# Explain: Why does increasing temperature often **increase** solubility for solids?

Answer 44

Higher temperatures **break** bonds, dissolving more solute. ## Footnote For gases, solubility decreases with increasing temperature because gas molecules escape more readily from the liquid.

Question 45

# Define: density

Answer 45

Measure of mass per unit volume of a substance. ## Footnote **Density** is a key physical property calculated as: Density = Mass ÷ Volume It is measured in units like g/cm³ or kg/m³.

Question 46

# Explain: Why does ice **float** on water?

Answer 46

Ice is **less dense** than liquid water due to its open hexagonal structure. ## Footnote As water freezes, its molecules arrange in a crystalline structure, increasing volume and decreasing density.

Question 47

# Define: principle of conservation of mass

Answer 47

Mass cannot be **created** or **destroyed** in a closed system. ## Footnote In chemical reactions, the total mass of reactants equals the total mass of products.

Question 48

# Explain: What is the **key idea** behind the **law of conservation of energy**?

Answer 48

Energy remains **constant** in a closed system, though it may change forms. ## Footnote This principle applies to mechanical, thermal, chemical, and other forms of energy.

Question 49

# Explain: What are the different **forms** **of energy**?

Answer 49

* Potential Energy * Chemical Energy * Gravitational Energy * Mechanical Energy * Nuclear Energy * Kinetic Energy * Radiant Energy * Thermal Energy * Motion Energy  * Sound Energy * Electrical Energy ## Footnote The principle of conservation of energy states that **energy cannot be created or destroyed**, **only transformed from one form to another**. All forms of energy adhere to this principle. **Potential Energy**: Stored energy due to position or state.   * **Chemical Energy**: Stored in the bonds of atoms and molecules (e.g., in food, batteries, and fuels).   * **Gravitational Energy**: Stored due to an object's height above the ground (e.g., a book on a shelf).   * **Mechanical Energy**: Stored in objects under tension or compression (e.g., a compressed spring).   * **Nuclear Energy**: Stored in the nucleus of an atom (e.g., in radioactive materials). **Kinetic Energy:** Energy of motion.   * **Radiant Energy:** Electromagnetic energy that travels in waves (e.g., light, X-rays, and radio waves).   * **Thermal Energy:** Energy of heat due to the movement of atoms and molecules. * **Motion Energy:** Energy of a moving object (e.g., a rolling ball).   * **Sound Energy:** Energy that travels as waves through a substance.   * **Electrical Energy:** Energy of moving electrons (e.g., in a wire).

Question 50

# Explain: What are **examples** of systems where mass is conserved?

Answer 50

* Chemical reactions * Physical changes like melting or boiling * Closed container experiments ## Footnote *Mass conservation* holds true in any system that does not exchange matter with its surroundings.

Question 51

# Explain: Significance of the **conservation of energy matter** in physics.

Answer 51

It **predicts** system behavior and ensures energy is accounted for in processes. ## Footnote This principle underpins laws like thermodynamics and mechanics.

Question 52

# Identify: The equation that **represents** the equivalence of mass and energy.

Answer 52

E = mc² ## Footnote Proposed by **Einstein**, this equation shows how mass can convert to energy and vice versa in nuclear reactions.

Question 53

# Identify: What happens to **total energy** in a closed system when **friction** is present?

Answer 53

It remains **constant** but may be transformed into heat. ## Footnote **Friction** dissipates mechanical energy as thermal energy, adhering to conservation laws.