Matter and Energy

I. Matter Matter is anything that has mass and volume Mass Volume Amount of matter Measured in grams (g) Volume Space matter occupies Measured in milliliters (mL), liters (L) or cubic centimeters (cm3 ) Two types of matter Pure substances Mixtures

A. Pure Substances Uniform composition Two Types The same throughout the sample Two Types Elements Compounds

1. Elements Simplest form of matter Smallest part called atom Cannot breakdown Represented using a capital letter or capital letter and lower case letter Carbon (C) Sodium (Na)

2. Compound Two or more elements chemically joined in a specific ratio Properties of the compound are different than the elements that make it up Compounds can be broken down Decomposed Examples Water (H2O) Hydrogen peroxide (H2O2)

B. Mixture Two or more substances physically joined in any ratio Keep the properties of their components (parts) Can be separated by physical means Exist in two forms Heterogeneous Homogeneous

Heterogeneous Homogeneous Visible difference between components (parts) No visible differences between components (parts) Called a solution Represented using (aq)

C. Properties of matter Physical Properties Chemical Properties Properties that can be observed without changing the substance Chemical Properties Properties that show how a substance reacts (changes)

Density—a physical property used to identify a substance Density = mass volume

Determine the density of an object with a mass of 59 Determine the density of an object with a mass of 59.3g and a volume of 84.2 mL. Density = Mass ÷ Volume d = 59.3g ÷ 84.2mL d = .704g/mL   Determine the mass of an object with a density of 0.94g/mL and a volume of 24.8 mL. Solve for mass Mass = Density x Volume m = 0.94g/mL x 24.8mL m = 23.3g Determine the volume of an object with a density of 1.47g/mL and a mass of 28.6g. Solve for volume Volume = Mass ÷ Density v = 28.6g ÷ 1.47g/mL v = 19.5mL

II. Energy Energy is the driving force behind change Cannot be created or destroyed It can change its form Two types of energy Kinetic Energy of motion Potential Stored energy

A. Measurements involving energy 1. Temperature Average kinetic energy of particles Measured using a thermometer (unit: degrees) Fahrenheit Celsius Kelvin To convert °F to °C -- use °C = 5/9( °F - 32) °C to ° F -- use °F = 9/5 °C + 32 °C to K -- use K = °C + 273 K to °C – use K = °C + 273

2. Calorimetry Measures the actual energy (q) in a system Related to mass (m), specific heat capacity (C) and temperature change (∆T) Measured using a calorimeter (unit: joules) To calculate energy use Reference Table T

How many joules are required to heat 40g water at 30°C to 80°C? q = m C ∆T q = 40g x 4.18J/g°C x 50°C q = 8360J

5000J were added to 30g water at 25°C. What is the new temperature? q = m C ∆T 5000J = 30g x 4.18J/g°C x ∆T 5000 = 125.4 x ∆T ∆T = 39.9 ~ 40 T new = 25 + 40 T new = 65°C

How many joules are needed to melt 100g ice at 0°C q = m Hfus q = 100g x 334J/g q = 33400J

III. Phases of Matter Solids Liquids Gases

A. Solids Matter that has specific shape and specific volume Atoms closely packed together Cannot be compressed

B. Liquids Matter that has a specific volume but takes the shape of the container Atoms are close but have some space between them Cannot be compressed Can be poured

C. Gases Matter that takes the shape and volume of the container Atoms have free space between them Compressible Can be poured

D. Phase Changes If energy is added… If energy is removed… Solid to liquid Melting Liquid to gas Boiling Solid to gas Sublimation Liquid to solid Freezing Gas to liquid Condensing Gas to solid Deposition

E. Phase Diagram   Boil Gas Temperature Melt Liquid Solid Time

  Heat of Vaporization Gas Temperature Heat of Fusion Liquid Solid Time

Kinetic Energy Potential Energy Temperature Potential Energy Kinetic   Kinetic Energy Potential Energy Temperature Potential Energy Kinetic Energy Kinetic Energy Time

F. Gas Laws Gas volume is controlled by pressure and temperature 1. Units Volume mL L cm3 Pressure atm kPa Temperature °C K is the only unit to be used

F. Gas Laws 2. Boyle Law As pressure increases, volume decreases P1V1 = P2V2 3. Charles Law As temperature increases, volume increases V1 = V2 T1 T2

Converting pressure units Convert 1.9 atm to kPa 101.3 kPa = x kPa 1 atm 1.9 atm (101.3)(1.9) = (1)(x) x = 192.7 kPa Convert 180 kPa to atm 101.3 kPa = 180 kPa 1 atm x atm (101.3) (x) = (1)(180) x = 1.8 atm

20 mL of CO2(g) at 1. 2 atm is compressed to 15 mL 20 mL of CO2(g) at 1.2 atm is compressed to 15 mL. What is the new pressure of the gas? P1V1 = P2V2 20 mL = volume (V) 1.2 atm = pressure (P) 15 mL = volume (V) (1.2 atm)(20 ml) = P2(15 ml) P2 = 1.6 atm

100 mL of NH3(g) at 25°C is heated to 50°C 100 mL of NH3(g) at 25°C is heated to 50°C. What is the new volume of the gas? V1 = V2 T1 T2 25°C + 273 = 298 K 50°C + 273 = 323 K 100 ml = V2 298 K 323 K (100)(323) = (298) V 32300 = 298V 298 298 V2 = 108 ml

When 500 mL of NH3(g) at 25°C was cooled, its new volume became 250 mL When 500 mL of NH3(g) at 25°C was cooled, its new volume became 250 mL. What is the Celsius temperature of the gas? V1 = V2 T1 T2 25°C + 273 = 298 K 500 ml = 250 mL 298 K T2 (298)(250) = (500) T 74500 = 500 T 500 500 T2 = 149 K 149 K = °C + 273 -124°C