Honors Chemistry Matter and Energy MATTER is anything that has mass and volume. - PowerPoint Presentation

Honors Chemistry Matter and Energy

MATTER is anything that has mass and volume. Do you know ALL the states of matter?

Kinetic Molecular Theory of Matter Assumptions: All matter consists of particles, such as, atoms, molecules, formula units. Particles are in constant motion (kinetic). These motions are vibrational (only for solids), translational, and/or rotational. Collisions are elastic.Theory explains the physical properties of matter.The state of matter of a substance depends upon the strength of attraction between particles. All 3 for liquids and gases.

Vibrational Combined Motions Translational Rotational

Gases Liquids Solids gas liquid solid little to no order most orderlow density high density (condensed states) compressible incompressible no definite volume or shape definite volume definite volume takes shape of container definite shape diffuses quickly slower diffusion (  with  temp) low diffusion low strength of attraction high strength of attraction

States of Matter melting freezing condensing evaporating deposition sublimation Solid Liquid Gas

Properties of Matter Physical = A characteristic of a substance that does not involve a chemical change Examples: texture, state of matter, density, hardness, boiling point

Properties of Matter Chemical = A property of matter that describes a substance’s ability to participate in chemical reactions. Examples: reacting with oxygen, light sensitivity

Physical Changes Do NOT change the identity Often change what the substance looks like Examples: mixing ice tea in water, crushing a rock, freezing water

Chemical Changes Alter the identity of the substance. The new substance has a different composition than the beginning substances. Examples: rusting and burning A shorthand way to express a chemical reaction is with a chemical equation.

Chemical Equations The substances on the left side of the arrow are called the reactants. They are the starting materials in the reaction. The substances on the right side of the arrow are called the products. They are the ending materials in a reaction. A + B C + D

Examples of Chemical Equations

Signs of a Chemical Change Gas production – bubbles, odor, fizz, smoke Color change Release or absorption of energy – light or temperature change Formation of a precipitate – a solid substance that falls out of solution

Classification of Matter

Classification of Matter Pure substances: A sample of matter with a definite composition; means definite chemical and physical properties. Includes: Elements and Compounds

Elements Made up of one type of atom. An atom is defined as the smallest unit of an element that maintains the properties of that element. Cannot be separated into simpler substances by chemical means.Represented by symbols.

Elements (cont.) Can exist as atoms or molecules. A molecule usually consists of two or more atoms. Ex. N 2 , O 2 , F2, Cl2, Br2, I2, H2, P4, S8Elements that have more than one form are called allotropes. Ex. Carbon (graphite and diamond)

Parts of the Periodic Table

Metal, Nonmetals and Metalloids (Semimetals): Metals Found on the LEFT side of the PT 2. Nonmetals Located on the RIGHT side of the PT 3. Metalloids - Properties of both metals & nonmetals - Good conductors of heat & electricity - High luster (shiny) - Ductile (can be drawn into thin wire) - Malleable (bends without breaking) - High melting points most solids at room temperature - High densities - Reacts with acids - Brittle (easy to break) - No luster (dull) - Neither ductile nor malleable - Nonreactive with acids - Insulators nonconductors (Semimetals)

Compounds Made up of 2 or more different elements combined in a fixed position. Can be separated through chemical means. Represented by formulas. Electrolysis allows chemists to distinguish between elements in compounds . Examples: CO2 and H2O

Elements Vs. Compounds

Mixtures: A combination of 2 or more substances that are not chem. combined. Heterogeneous Mixture: Composed of dissimilar components; Can see the parts A.K.A. Mechanical Mixture Ex. Cookie, salad, asphalt Homogenous Mixture: Uniform structure or composition throughout A.K.A. SolutionEx. Lemonade, steel, airAlloy: A solid homogeneous mixture (14 caret gold, steel, pewter)

Examples of Alloys Brass is an alloy of copper and zinc. Steel is an alloy of carbon and iron. Bronze is an alloy of copper and tin.

Microscopic look at mixtures

Counting Atoms in Compound Step 1: List all elements present Step 2: Identify the coefficient Step 3: Count the number of atoms of each element in the compound. Step 4: Multiply the coefficient by the subscript Step 5: Add up all the atoms

Counting Atoms Na 2 SO 4 Ca(OH) 2 3 Fe2(SO3)3

Separating Heterogeneous Mixtures Filtration: Pour liquid through paper and collect residue (solid)

Separation of Homogeneous Mixtures Distillation: Separation based on a difference in boiling points

Another Look at Distillation Distillation Demo A Closer Look at Distillation

Separation of Homogeneous Mixtures Crystallization: Evaporate liquid and solid will recrystallize

Separation of Homogeneous Mixtures Chromatography: Separation of pigments of dye

Percent Concentration of Solutions A measure of the amount of solute in a solution. % Concentration = mass of solute x 100 mass of solutionNote: Solute + Solvent = Solution

Solution Definition: a homogeneous mixture of 2 or more substances in a single physical state Parts: solute and solvent (usually water) Types: Physical states: solid (alloys), liquid, gas Miscible vs. Immiscible Miscible: Liquids that dissolve freely in one another in any proportionImmiscible :Liquid solutes and solvents that are not solubleSaturated, Unsaturated and SupersaturatedDilute vs. Concentrated

Saturated – soln containing the max amt of solute Unsaturated – soln containing less solute than a sat soln under the existing conditions Supersaturated – contains more dissolved solute than a saturated solution under the same conditions Solubility Curves

supersaturated solution(stirred) Supersaturated Solution of Sodium Thiosulfate

Solubility (physical change) Definition: mass of solute needed to make a saturated solution at a given temperature solution equilibrium in a closed system dissolution ↔ crystallization Unit = g solute/100 g H2O

Solubility of solids in liquids For most solids, increasing temperature, increases solubility. In general, “like dissolves like”. Depends on Type of bonding Polarity of molecule Intermolecular forces between solute and solvent

Saturated sol’n Supersaturated solution Unsaturated solution At 20 o C, a saturated solution contains how many grams of NaNO 3 in 100 g of water? What is the solubility at 70 o C? 135 g/100 g water What kind of solution is formed when 90 g NaNO 3 is dissolved in 100 g water at 30 o C? unsaturated What kind of solution is formed when 120 g NaNO 3 is dissolved in 100 g water at 40 o C? supersaturated 90 g

Solubility of Gases Gases are less soluble at high temperatures than at low temperatures Increasing temperature, decreases solubility. Increasing pressure, increases solubility.

Increasing pressure, increases solubility. The quantity of gas that dissolves in a certain volume of liquid is directly proportional to the pressure of the gas (above the solution).

Effervescence – rapid escape of gas dissolved in liquid

Factors Affecting Solubility Increase surface area of solute (crushing) Stir/shake Increase temperature

Energy and Change Energy is the capacity to do work . All physical and chemical changes require energy . Endothermic - describes a process in which heat is absorbed from the environment .Exothermic – describes a process in which heat is released into the environment.

Law of Conservation of Energy Energy is neither created, nor destroyed. It just changes forms.

Types of Energy Potential energy – stored energy Kinetic energy – energy of motion

Heat Transfer Transfer of heat may not affect temperature. During a phase change, the temperature will remain constant until all of the substance has changed state. The temperature will increase when a substance is a solid, liquid, or gas.

Kinetic Theory of Matter Gases posses the greatest amount of kinetic energy. Two factors that determine the state of matter of a substance: speed of particles and distance There are two factors contribute to the attraction between the particles.

Kinetic Theory of Matter Substances change phases when they overcome these attractions. The overall kinetic energy will not change until the entire substance has completely changed. Comparison of the three states of matter

Kinetic-Molecular Theory and Gases 1. Gases are small particles that have mass . These particles are usually molecules, except for the Noble Gases.

2. The particles in gases are separated by relatively large distances.

3. The particles in gases are in constant rapid motion (random). 4. Gases exert pressure because their particles frequently collide with the walls of their container and each other.

5. Collisions of gas particles are perfectly elastic . Inelastic Collision Elastic Collision

Gas particles do not slow down when hitting each other or the walls of their container.

6. Temperature of a gas is simply a measure of the average kinetic energy of the gas particles. High temp. = high KE, Low temp. = low KE

7. Gas particles exert no force on one another. Attractive forces are so weak between particles they are assumed to be zero .

Boyle’s Law Pressure - Volume Relationship . The pressure & volume of a sample of gas at constant temperature are inversely proportional to each other.Indirect P1 V 1 = P 2 V 2

Boyle’s Law

Boyle’s Law Problem A gas has a volume of 300. mL under a pressure of 740. mm of mercury. If the temperature remains constant, calculate the volume when under a pressure of 750. mm Hg. P 1V1 = P2V2

Charles’ Law : Temperature - Volume Relationship . At constant pressure the volume of a fixed amount of gas is directly proportional to its absolute temperature. Law assumes n is constant.Direct *Temperatures must be in Kelvin! K = ° C + 273

Balloon in cool and cold water:

Charles’s Law

A gas sample at 83ºC occupied a volume of 1470 m3 . At what temperature, in ºC, will it occupy a volume of 1250 m 3 ? V 1 = 1470 m3 V2 = 1250 m3T1 = 83°C = 356 K T2 = ?T 2 = 30.°C Charles’ Law Problem

The pressure of a fixed volume of gas is directly proportional to its absolute temperature . Law assumes n is constant. Direct P1 = P2 T1 T 2 *Temperatures must be in Kelvin! K = ° C + 273 Gay-Lussac’s Law

Gay-Lussac’s Law

Before a trip, the pressure in a car tire was 1.80 atm at 21 oC. At the end of the trip, the pressure gauge reads 1.90 atm. Calculate the temperature, in Celsius, of the air inside the tire at the end of the trip. Assume the tire volume does not change. P 1 = 1.80 atm P2 = 1.90 atmT1 = 21°C = 294 K T2 = ? T 2 = 37°C Gay-Lussac’s Law

The Combined Gas Law “ Choyles ” This law can be used to determine how changing two variables at a time affects a third variable.

A gas occupies 72.0 mL at 25 ° C and 198 kPa . Convert these to standard conditions. What is the new volume? P1 = 198 kPa P2 = 101.325 kPa V1 = 72.0 mL V 2 = ? T1 = 298 K T 2 = 273 K

Gases in a mixture behave independently of each other. The total pressure of a gaseous mixture equals the sum of the partial pressures of the individual gases in a mixture. Partial pressure = individual pressure of a gas in a mixturePT = p1 + p2 + p3 + …Dalton’s Law of Partial Pressures

Dalton’s Law of Partial Pressures : Example #1) A flask contains a mixture of oxygen, argon, and carbon dioxide with partial pressures of 745 torr , 0.278 atm , and 391 torr respectively. What is the total pressure in the flask? P T = P a + P b + P c + …

In the lab, gases are collected over water (water displacement). As a result, water vapor contributes to the total pressure.P T = p dry gas + p water vapor where p water vapor varies with temperatureDalton’s Law of Partial Pressures

T ( o C ) P (mm Hg) T ( o C) P (mm Hg) T ( o C) P (mm Hg) T ( o C) P (mm Hg) 0 4.6 26 25.2 51 97.2 76 301.4 1 4.9 27 26.7 52 102.1 77 314.1 2 5.3 28 28.4 53 107.2 78 327.3 3 5.7 29 30.0 54 112.5 79 341.0 4 6.1 30 31.8 55 118.0 80 355.1 5 6.5 31 33.7 56 123.8 81 369.7 6 7.0 32 35.7 57 129.8 82 384.9 7 7.5 33 37.7 58 136.1 83 400.6 8 8.1 34 39.9 59 142.6 84 416.8 9 8.6 35 42.2 60 149.4 85 433.6 10 9.2 36 44.6 61 156.4 86 450.9 11 9.8 37 47.1 62 163.8 87 468.7 12 10.5 38 49.7 63 171.4 88 487.1 13 11.2 39 52.4 64 179.3 89 506.1 14 12.0 40 55.3 65 187.5 90 525.8 15 12.8 41 58.3 66 196.1 91 546.1 16 13.6 42 61.5 67 205.0 92 567.0 17 14.5 43 64.8 68 214.2 93 588.6 18 15.5 44 68.3 69 223.7 94 611.0 19 16.5 45 71.9 70 233.7 95 634.0 20 17.5 46 75.7 71 243.9 96 658.0 21 18.7 47 79.6 72 254.6 97 682.0 22 19.8 48 83.7 73 265.7 98 707.3 23 21.1 49 88.0 74 277.2 99 733.2 24 22.4 50 92.5 75 289.1 100 760.0 25 23.8  

Eudiometer Piece of glassware used to measure the change in volume of a gas . It is similar to a graduated cylinder. It is closed at the top end with the bottom end immersed in water or mercury. The liquid traps a sample of gas in the cylinder, and the graduation allows the volume of the gas to be measured.

Example #2) Atmospheric pressure is 101.3kPa, and air is a mixture of N 2 , O 2, and Ar as 78.0%, 21.0%, and 1.0%, respectively. Calculate the partial pressure of O 2 . 21.3 kPaExample #3) Hydrogen gas is collected by water displacement at 18°C. Air pressure on that day is 744.0 mm. Calculate the pressure due to the dry hydrogen gas. 728.5 mm Hg