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CHAPTER 3: MATTER & ENERGY

Problems: 1-50, 55-70, 103-105, 107-111, 113 3.2 What is Matter? Matter: Anything that has mass and occupies volume We study matter at different levels:

macroscopic: the level that can be observed with the naked eye – e.g. geologists study rocks and stone at the macroscopic level microscopic: the level that can be observed with a microscope – e.g. scientists study tiny animals, plants, or crystals at microscopic level particulate: at the level of atoms and molecules, also called atomic or molecular level

– cannot be observed directly even with the most powerful microscopes – where the term “nanotechnology” comes from since many atoms and

molecules are about a few nanometers in size

Substances like water can be represented using different symbols (e.g H2O) and models.

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3.3 Classifying Matter According to Its State: Solid, Liquid, and Gas Matter exists in one of three physical states: solid, liquid, gas

solid: Has definite shape and a fixed (or constant) and rigid volume – Particles only vibrate in place. liquid: Has a fixed (or constant) volume, but its shape can change. – Takes the shape of its container because particles are moving – Particles are packed closely together but can move around each other. gas: Volume is variable, and particles are far apart from one another. – Takes the shape of the container because particles are moving → If container volume expands, particles move apart to fill container. → If container volume decreases, particles move closer together. → Gases are compressible—i.e., can be forced to occupy a smaller volume. – Particles are in constant random motion.

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3.4 Classifying Matter According to Its Composition: Elements, Compounds, and Mixtures We can classify matter into pure substances and mixtures: pure substance: a single chemical, consisting of only one kind of matter – There are two types of pure substances: elements and compounds. – In the figure below, copper rods are an example of an element, and sugar is an example

of a compound. mixture: consists of two or more elements and/or compounds – Mixtures can be homogeneous or heterogeneous: – Homogeneous mixtures have a uniform appearance and composition because the

particles in them mix uniformly (e.g. solutions like sweetened tea below) – Heterogeneous mixtures do not have a uniform composition.

– e.g. chocolate chip cookie, water and C8H18 mixture below shown as separate layers

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elements: – consist of only one type of atom – atoms cannot be broken down into smaller

components by chemical reaction – e.g. copper wire (Cu), sulfur powder (S8) – Examples also include sodium (Na),

barium (Ba), hydrogen gas (H2), oxygen gas (O2), and chlorine gas (H2).

compounds: – consist of more than one type of atom and

have a specific chemical formula – Examples include hydrogen chloride (HCl),

water (H2O), sodium chloride (NaCl) which is table salt, barium chloride (BaCl2)

Two or more pure substances combine to form mixtures. mixtures: – consist of many compounds and/or elements,

with no specific formula – Matter having variable composition with definite or varying properties – can be separated into component

elements and/or compounds – e.g., any alloy like brass, steel, 10K to

18K gold; sea water, carbonated soda; air consists of nitrogen, oxygen, and other trace gases.

The image at the right shows that air is a mixture of mostly nitrogen (N2 in blue) and some oxygen (O2 in red) while salt water consists of salt (Na+ and Cl- ions or charged particles) dissolved in water. Example: Is salt water a homogeneous or heterogeneous mixture? Explain.

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3.5 How We Tell Different Kinds of Matter Apart: Physical and Chemical Properties and Changes

The characteristics that distinguish one substance from another are called properties.

Physical Properties: inherent characteristics of a substance independent of other substances

– physical state (solid, liquid, gas) – electrical & heat conductivity – color – odor – density – hardness – melting and boiling points – solubility (does/does not dissolve in water)

Chemical Properties: how a substance reacts with other substances – e.g. hydrogen reacts explosively with oxygen 3.6 How Matter Changes: Physical and Chemical Changes physical change: – a process that does not alter the chemical makeup of the starting materials – Note in the figure below that the H2O molecules remain H2O regardless of the physical

state (solid, liquid, or gas). → Changes in physical state are physical changes.

– Other examples of physical changes include hammering gold into foil, dry ice subliming – Dissolving table salt or sugar in water is also a physical change. – A substance dissolved in water is the fourth physical state, aqueous.

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Know the terms for transitions from one physical state to another:

freezing: liquid → solid condensing: gas → liquid melting: solid → liquid evaporating (or vaporizing): liquid → gas

Two less common transitions: sublimation: solid → gas (e.g. dry ice sublimes) deposition: gas → solid (e.g. water vapor deposits on an icebox) chemical change: – a process that does change the chemical makeup of the starting materials

– We can show H2 and O2 reacting to form water (H2O) below. Since the H2O has a different chemical makeup than H2 and O2, this is a chemical change.

– Other examples of chemical changes: – e.g. oxidation of matter (burning or rusting), release of gas bubbles (fizzing) ,

mixing two solutions to form an insoluble solid (precipitation), and other evidence indicating the starting materials (reactants) were changed to a different substance.

– The following examples are all chemical changes that convert the reactants to completely different compounds and/or elements.

release of gas bubbles

(fizzing)

formation of insoluble solid

(precipitation)

oxidation (burning or rusting)

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Example 1: Consider the following molecular-level representations of different substances:

For each figure above, indicate if it represents an element, a compound, or a mixture AND if it represents a solid, liquid, or gas. A: element compound mixture solid liquid gas

B: element compound mixture solid liquid gas

C: element compound mixture solid liquid gas

D: element compound mixture solid liquid gas

E: element compound mixture solid liquid gas

F: element compound mixture solid liquid gas

Ex. 2: Circle all of the following that are chemical changes:

burning condensing dissolving rusting vaporizing precipitating

A B

D E

C

F

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Ex. 3: Classify the following as pure substances or mixtures: A homogeneous liquid whose temperature stays constant while boiling.

Pure substance

Mixture

Granite—a rock with several visible minerals in it. Pure substance

Mixture

A red solid that turns blue when heated and releases water that is always 30% of the solid’s mass.

Pure substance

Mixture

A gas that when cooled and compressed, a liquid condenses out but some gas remains.

Pure substance

Mixture

Ex. 3: Classify the following as pure elements or compounds: Chlorine, Cl2 Element Compound

Table sugar, C12H22O11 Element Compound

A red solid that turns blue when heated and releases water that is always 30% of the solid’s mass.

Element Compound

A brown-red liquid that, when energy is applied to it in any form, causes only physical changes in the material, not chemical.

Element Compound

3.7 Conservation of Mass: There is No New Matter

Chemical Reaction: REACTANTS → PRODUCTS (starting materials) (substances after reaction)

For the reaction: C + O2 → CO2 The reactants are carbon and oxygen gas, and the product is carbon dioxide.

Antoine Lavoisier (1743-1794), a French chemist, carried out experiments on combustion by burning different substances and measuring their masses before and after burning. – He found that there was no change in the overall mass of the sample and air around it. → Law of Conservation of Mass: Matter is neither created nor destroyed in a chemical reaction, so mass is conserved.

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Mass of the product(s) in a reaction must be equal to the mass of the reactant(s). For example: 11.2 g hydrogen + 88.8 g oxygen = 100.0 g water

Ex. 1: Methane burns by reacting with oxygen present in air to produce steam and

carbon dioxide gas. Calculate the mass of oxygen that reacts if burning 50.0 g of methane produces 112.3 g of steam and 137.1 g of carbon dioxide.

3.8 Energy: the capacity to do work potential energy (PE): energy due to position or its composition (chemical bonds) – A 10-lb bowling ball has higher PE when it is 10 feet

off the ground compared to 10 inches off the ground → Greater damage on your foot after falling 10 feet

compared to falling only 10 inches – In terms of chemical bonds, the stronger the bond, → more energy is required to break the bond, → the higher the potential energy of the bond

kinetic energy (KE): energy associated with an object’s motion – e.g. a car moving at 75 mph has much greater KE than the same car moving at 15 mph → Greater damage if the car crashes at 75 mph than at 15 mph Six Forms of Energy: heat, light, chemical, electrical, mechanical, and nuclear – Each can be converted to another. Example: Identify at least two types of energy involved for each of the following: 1. When you turn on a lamp 2. When using solar panels 3. At the Springfield Power Plant in The Simpsons

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Energy changes accompany physical and chemical changes due to changes in potential and kinetic energy.

Kinetic Energy and Physical States

Solids have the lowest KE of the three physical states – Highest attraction between particles → particles are fixed Liquids have slightly higher KE than solids – Particles are still attracted to each other but can move past one another → particles are less restricted Gases have greatest KE compared to solids and liquids – Attractive forces completely overcome, so particles fly freely within container → particles are completely unrestricted

3.10 Temperature: Random Molecular and Atomic Motion Heat: Energy that is transferred from a body at a higher temperature to one at a lower

temperature → heat always transfers from the hotter to the cooler object! – "heat flow" means heat transfer

Heat Transfer and Temperature – One becomes hotter by gaining heat. – One becomes colder by losing heat—i.e., when you “feel cold”, you are actually losing heat! Ex. 1: Fill in the blanks to indicate how heat is transferred:

a. You burn your hand on a hot frying pan. ________________ loses heat, and ______________ gains the heat. b. Your tongue feels cold when you eat ice cream. ________________ loses heat, and ______________ gains the heat.

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Ex. 2: A small chunk of gold is heated in beaker #1, which contains boiling water. The gold chunk is then transferred to beaker #2, which contains room-temperature water.

a. The temperature of the water in beaker #2 _____. ↑ ↓ stays the same

b. Fill in the blanks: _____________ loses heat, and ____________ gains the heat. 3.9 Energy and Chemical and Physical Change endothermic change: a physical or chemical change that requires energy or heat to occur

– boiling water requires energy: H2O(l) + heat energy → H2O(g)

– electrolysis of water requires energy: 2 H2O(l) + electrical energy → 2 H2(g) + O2(g) exothermic change: a physical or chemical change that releases energy or heat

– water condensing releases energy: H2O(g) → H2O(l) + heat energy

– hydrogen burning releases energy: 2 H2(g) + O2(g) → 2 H2O(g) + heat energy For physical changes, consider whether the reactants or products have more kinetic energy. – If the reactants have greater kinetic energy than the products → exothermic process. – If the products have greater kinetic energy than the reactants → endothermic process.

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3.10 Temperature The Kelvin (K) scale gives a direct measure of the thermal energy present in a sample. Many countries, and most scientists use Celsius (or Centigrade) (°C) scale, which is set by the freezing point and boiling point of water. The United States persist in using the Farenheit (°F) scale, which is set using mercury (the only metal that exists as a liquid at room temperature) as a reference.

The three temperature scales are related according to the following equations: °F = (9/5) °C + 32 °C = (5/9)(°F - 32) K = °C + 273.15

Ex. 1: The lowest temperature measured on Earth is -128.6°, recorded at Vostok, Antarctica, in July 1983. What is this temperature on the Celsius and Kelvin scales?