unit 2: nature of matter and kinetic theory. part 1: the nature of matter

Unit 2: Nature of Matter and Kinetic Theory

Part 1: The Nature of Matter

properties = characteristics and behavior of matter (includes changes that matter undergoes).

What color is it?Is it solid, liquid or gasIs it reactive?

structure = composition

• what matter is made of

• how matter is organized.

How do we classify matter?

• Examples of physical properties :

• solubility, - dissolves in water?• melting point, boiling point

• color,• density,

• electrical conductivity,

• physical state (solid, liquid, or gas).

• physical change - change in matter that does not involve a change in the chemical identity

• Change of state is a physical change:

Classify by purityIs it a pure substance or mixture?

Pure substance = sample of matter that has definite chemical and physical properties, can be either an element or a compound

Classifying Matter

compound = pure substance that can be broken down into simpler substances.

element = substance that cannot be broken down into simpler substances.

Element or Compound?

salt

gold

Compounds Are More Than One Element

formula = combination of the chemical symbols that show what elements make up a compound and the number of atoms of each element

Compound Formula

caffeine

salt

water

C8H10N4O2

NaCl

H2O

Compounds Are More Than One Element

****The properties of the compound are different from the properties of the elements that compose the compound.

silver + bromine = silver bromide

substance is not changed = no fixed composition

the basic identity of each

Mixture = made up of different kinds of matter

Pure Substance or Mixture?

•Homogeneous mixtures are the same throughout.

• Also known as a solution.

Pure substance or a mixture?

• When you dissolve sugar in water, sugar is the solute—the substance being dissolved.

• The substance that dissolves the solute is the solvent. in this case it is water

solute + solvent = solution

• When the solvent is water, the solution is called an aqueous solution.

heterogeneous mixture is one with different compositions, depending upon where you look

Pure Substance or Mixture?

Pure Substance

Mixture

element

compound

homogenous

heterogeneous

Matter

• a substance must be separated chemically

• a mixture can be separated physically

An example of a pure substance in everyday life is _____.

a. pond water

b. a cola drink

c. sugar

d. concrete

c. sugar

A soft drink is an example of a(n) _____.

a. compound

b. heterogeneous mixture

c. element

d. homogeneous mixture

d. homogenous mixture

Identify each of the following as either a compound or a mixture.

A. sand

B. water

C. juice

mixture

compound

mixture

In ocean water, salt is a(n) _____.

a. alloy

b. solution

c. solute

d. solvent

c. solute

pure substance?element or compound?

a mixture?Heterogeneous or

homogenous Aluminum foil

Pure substance, element

pure substance?element or compound?

a mixture?Heterogeneous or homogenous

• bowl of cereal

mixture, heterogeneous

pure substance?element or compound?

a mixture?Heterogeneous or homogenous

• whipped cream

Mixture, homogenous

pure substance?element or compound?

a mixture?Heterogeneous or homogenous

• oil and vinegar dressing

Mixture, heterogeneous

pure substance?element or compound?

a mixture?Heterogeneous or homogenous

• aspirin - acetylsalicylic acid

Pure substance, compound

pure substance?element or compound?

a mixture?Heterogeneous or homogenous

• orange juice with pulp

Mixture, heterogeneous

pure substance?element or compound?

a mixture?Heterogeneous or homogenous

• gold

Pure substance, element

pure substance?element or compound?

a mixture?Heterogeneous or homogenous

• salt

Pure substance, compound

pure substance?element or compound?

a mixture?Heterogeneous or homogenous

• peanut butter

Mixture, homogenous

Chemical Properties• Chemical properties are those that can

be observed only when there is a change in the composition of the substance.

• Rusting is a chemical reaction in which iron combines with oxygen to form a new substance, iron oxide.

Examples of chemical property:

• flammability

• reactivity

Chemical Changes

chemical change - the change of one or more substances into other substances.

• A chemical property always relates to a

chemical change = chemical reaction.

- production of bubbles- release or absorption of energy

- color change

***only way to be sure is to check the composition of the sample before and after the change.

Clues that a chemical change has occurred:

Below are listed changes that can be observed in everyday life. Tell whether it is a physical change or a chemical change.

1.an icicle melting

2.charcoal burning

3.magnetizing a piece of steel

4.iron rusting

5.rubbing alcohol evaporating from the skin

physical change

chemical change

physical change

chemical change

physical change

chemical change involves only a rearrangement of the atoms. Atoms DO NOT just appear or disappear.

******Law of Conservation of Mass****** In a chemical change, matter is neither created nor destroyed.

Chemical Reactions

Chemical Reactions and Energy

• All chemical changes also involve some sort of energy change.

• Energy is either taken in or given off as the chemical change takes place. Energy is the capacity to do work.

• Work is done whenever something is moved.

Chemical Reactions and Energy

• Energy is also produced and released in the form of heat and light.

Chemical reactions that GIVE OFF heat energy are called exothermic reactions.

• Chemical reactions that ABSORB heat energy are called endothermic reactions.

Chemical Reactions and Energy

Classify each of the following as a chemical or physical property.

A. density

B. reactivity

C. color

D. melting point

physical property

chemical property

physical property

physical property

Part 2: The Kinetic Theory

• States of Matter–solid–liquid–gas–plasma

Intermolecular Forces (IMF)

• Attractive forces between molecules.

Much weaker than chemical bonds within molecules.

The Kinetic Theory of Matter

1. Matter is composed of PARTICLES.

2. Particle movement is rapid, constant, and random (Brownian motion)

The Kinetic Theory of Matter

3.All collisions are perfectly ELASTIC (NO energy lost).

Kinetic theory of matter

Kinetic energy (K.E.) = energy of motion

• gases have the least restriction on motion– have the most K.E.

• solids have the most restriction on motion– have the least K.E.

Kinetic model of gases

• Gases: matter with variable shape and variable volume

• Gas particles move in a straight line until they collide with container or each other

Kinetic model of liquids• Liquids: matter with variable

shape and definite volume

• Particles slide past each other but are so close together they do not move in a straight line

Kinetic model of solids

• solids: matter with definite shape and definite volume

• Particles cannot move past each other, they are in constant motion bouncing off neighbors

Other forms of matter• Plasmas - gaseous mixture of ions

-exists at high temperatures

• most common form of matter in the universe but least common on Earth itself

Plasmas continued

• an ionized gas that conducts electricity -forms at very high temps when matter absorbs energy and breaks apart

• The sun is made of plasma- also found in fluorescent lights

Temperature and kinetic energy

• temperature—the measure of the average K.E. of particles in a sample

• Kelvin (K) – SI base unit of temperature; measures average K.E.

Temperature and kinetic energy

• When temp increases, particle motion increases.

• When temp decreases, particle motion decreases.

A temp of 300 K has twice the kinetic energy as 150 K.

Temperature and kinetic energy

• 0 Kelvin = absolute zero = no molecular motion

• No degrees sign ( ° ) is used with Kelvin numbers

• There will never be negative numbers for Kelvin temperatures!.

density compressibility

intermolecular forces

solid most dense difficult to compress

strong

liquid

gas least dense easily compressed

weak

Comparing solids, liquids, and gases

Kinetic energy

space between particles

organization

solid least amount of kinetic energy

very little space between particles

most organized

liquid

gas most amount of kinetic energy

a lot of space between particles

least organized

Comparing solids, liquids, and gases

Changing states and energy changes

• Going from a more energetic state (gas) to a less energetic state (solid) requires a release of energy–exothermic

• Going from less energetic (solid) to more energetic (gas) requires absorption of energy -- endothermic

Vapor Pressure and boiling

• Vapor PressureVapor Pressure - pressure of vapor above a liquid at equilibrium

•high vapor pressure = volatile•volatile = easily evaporates

•The greater the fraction of molecules which can escape the liquid, the greater the vapor pressure

• What happens to the vapor pressure if you increase the temperature of a liquid in a closed container?

–causes the vapor pressure above the liquid to increase.

equilibrium vapor pressure - when the number of vapor molecules rejoining the water equals the number leaving to go into the vapor phase

• If there is equilibrium between the liquid state and the gas state, what is true about the rate of evaporation and the rate of condensation?

• They are equal

Vapor pressure and boiling point

• Boiling Point - temp at which v.p. of liquid equals external pressure

-depends on atmospheric

pressure & IMF

Normal B.P. - b.p. at 1 atm

Effects of Intermolecular Forces (IMF)

• When IMF’s are weak–vapor pressure is high–volatility is high–boiling point is low

Heat of Fusion

• Melting point – temp of a solid when it becomes a liquid= freezing point (temp when liquid

becomes a solid)

B. Heating Curves

Freezing/Melting point

Solid

Liquid

Boiling point

Gas

Heating Curves• IMPORTANT: temp does not change

during the actual phase change.• Increasing the temp will only make

the change happen faster.

Phase Diagrams• Shows the phases of a

substance at different temps and pressures.

triple point -the point (temperature and pressure) on a phase diagram at which three phases of a substance can coexist.

All six phase changes can occur at the triple point: freezing and melting, evaporation and condensation, sublimation and deposition.

Phase Diagrams

critical point -at extremely high temperatures and pressures, the liquid and gaseous phases become indistinguishable, in what is known as a supercritical fluid

The End!