CHAPTER 3: MATTER & ENERGY
Problems: 1-50, 57-58, 61-74, 107, 109-115
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.
CHEM 121: Tro Chapter 3 v0916 page 1 of 14 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.
CHEM 121: Tro Chapter 3 v0916 page 2 of 14 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
CHEM 121: Tro Chapter 3 v0916 page 3 of 14 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.
CHEM 121: Tro Chapter 3 v0916 page 4 of 14 3.5 How We Tell Different Kinds of Matter Apart: Physical and Chemical Properties
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.
CHEM 121: Tro Chapter 3 v0916 page 5 of 14 Know the terms for transitions from one physical state to another:
freezing:
melting:
liquid → solid
solid → liquid
condensing: gas → liquid
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)
3.6 How Matter Changes: Physical and Chemical Changes
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)
CHEM 121: Tro Chapter 3 v0916 page 6 of 14 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
CHEM 121: Tro Chapter 3 v0916 dissolving
rusting
vaporizing
precipitating
page 7 of 14 3.7 Conservation of Mass: There is No New Matter
Chemical Reaction:
REACTANTS
→
(starting materials)
PRODUCTS
(substances after reaction)
Consider the reaction below:
2 H2(g)
+
→
O2(g)
2 H2O(g)
The reactants are hydrogen (H2) gas and oxygen (O2) gas, and the product is steam.
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 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.
– Since the atoms are simply rearranged (not
created or destroyed) in a chemical reaction,
the total mass of the products must always
equal the total mass of the reactants.
Antoine Lavoisier with his wife and collaborator,
Marie-Anne-Pierrette Paulze
http://www.metmuseum.org/art/collection/search/436106
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.
CHEM 121: Tro Chapter 3 v0916 page 8 of 14 3.8 Energy: the ability to do work or produce heat
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 55 mph has much greater KE than the same car moving at 15 mph
→ Greater damage if the car crashes at 55 mph than at 15 mph
– Consider the Wile E. Coyote examples of PE and KE.
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:
a. When you turn on a lamp: ____________________________________
b. When using solar panels: ____________________________________
c. At the Springfield Power Plant in The Simpsons: _______________________________
http://cdn.appstorm.net/iphone.appstorm.net/files/2012/08/tapped_1.jpg
CHEM 121: Tro Chapter 3 v0916 page 9 of 14 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
Temperature:
A measure of the average kinetic energies of the particles in a substance
– i.e., a measure of the random motions of the particles in a substance
Heat: A measure of the total energy of the particles in a system (also called thermal energy)
Thermal energy is the kinetic energy associated with the motion of particles.
– Proportional to a substance’s temperature
– Increases with the size of a sample
Example: Consider the two beakers at the right
which both contain boiling water (at 100°C).
1. Which beaker has water molecules with
higher average kinetic energy?
(a)
(b)
neither
2. Which beaker contains water with higher thermal energy?
CHEM 121: Tro Chapter 3 v0916 (a)
(b)
neither
page 10 of 14 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
– View “Nova: Making Stuff Colder”: https://www.youtube.com/watch?v=0evHKYwI90M&t=0m39s
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.
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.
Ex. 3: Why do surfaces like a stone countertop or a metal railing inside a building feel cold?
Aren’t they at room temperature like you? Explain.
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
– electrolysis of water requires energy:
→ H2O(g)
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)
– hydrogen burning releases energy:
2 H2(g) + O2(g) → 2 H2O(g) + heat energy
CHEM 121: Tro Chapter 3 v0916 →
H2O(l)
+
heat energy
page 11 of 14 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.
system: that part of the universe being studied
surroundings: the rest of the universe outside the system
For chemical changes, observe if the surroundings (including you) feel hotter or colder after the
reaction has occurred.
– If the surroundings are hotter, the reaction released heat → exothermic reaction.
– If the surroundings are colder, the reaction absorbed heat → endothermic reaction.
Ex. 1: Circle all of the following changes that are exothermic:
freezing
CHEM 121: Tro Chapter 3 v0916 vaporizing
sublimation
melting
deposition
page 12 of 14 Ex. 2: A student adds ammonium chloride (NH4Cl) salt to a test tube containing water and notices
that the test tube feels colder as the ammonium chloride dissolves. Is this process
exothermic or endothermic? Explain.
Ex. 3: A student mixes two solutions, hydrochloric acid and sodium hydroxide, and notices the
beaker with the substances feels hotter as they mix. Is this reaction exothermic or
endothermic? Explain.
Units of Energy
calorie (cal): unit of energy used most often in the US
– amount of energy required to raise the temperature of 1 g of water by 1˚C
1 cal ≡ 4.184 J
(Note: This is EXACT!)
– But a nutritional calorie (abbreviated Cal) is actually 1000 cal:
1 Cal = 1 kcal = 4.184 kJ
joule (J): SI unit of energy
– To recognize the size of a joule, note that
1 watt = 1
→ So a 100-watt light bulb uses 100 J every second.
– Heat is also often reported in kilojoules (kJ), where 1 kJ = 1000 J
Electricity usage on our electricity bills are generally reported in kilowatt-hour (kWh).
Example: A 100 W lightbulb running for 10.0 hours requires 1 kWh of energy. If the average
Seattle home uses 25 kWh/day, this is equal to how many 100 W light bulbs running
nonstop for one day?
CHEM 121: Tro Chapter 3 v0916 page 13 of 14 Energy and Food Values
food value: The amount of heat released when food is burned completely, usually reported in
kJ/g food or Cal/g food.
– Most of the energy needed by our bodies comes from carbohydrates and fats, and the
carbohydrates decompose in the intestines into glucose, C6H12O6.
– The combustion of glucose produces energy that is quickly supplied to the body:
C6H12O6(g) +
6 O2(g) →
6 CO2(g) + 6 H2O(g) + heat energy
– The body also produces energy from proteins and fats, which can be stored because fats are
insoluble in water and produce more energy than proteins and carbohydrates.
– The energy content reported on food labels is generally determined using a bomb calorimeter
in which food is burned, and the change in temperature provides the total energy released.
3.11 Temperature Changes: Heat Capacity
heat capacity: amount of heat necessary to raise the temperature of a given amount of any
substance by 1°C; in units of J/°C
specific heat capacity: amount of heat necessary to raise the temperature of 1 gram of any (or
specific heat)
substance by 1°C; has units of J/g°C
Ex. 1: Consider water’s specific heat (4.184 J/g·°C) and the specific heats of rocks and other
solids (1.3 J/g·°C for dry Earth, 0.88 J/g·°C for concrete, 0.49 J/g·°C for steel). If samples of
each had the same initial temperature and were then heated with the same amount of
energy. Which substance would have the following?
highest final temperature: ____________ lowest final temperature: ____________
Thus, because water covers most of the Earth, water can absorb a lot more energy before its
temperature starts to rise.
→ Water helps to regulate temperatures on Earth within a comfortable range for humans.
→ Why coastal regions have less extreme temperatures compared to desert regions
CHEM 121: Tro Chapter 3 v0916 page 14 of 14