States of Matter

Chapter 3

Matter and Energy

States of Matter (cont.) Gases

1 Liquids Properties

• Characteristics of the substance under observation • Properties are: – Directly observable – The way something interacts with other substances in the universe

Universe Classified Properties of Matter

• Matter: the part of the universe that has • Physical Properties: the characteristics of mass and volume matter that can be changed without changing its composition • Chemistry is the study of matter – Characteristics that are directly observable – The properties of different types of matter • Chemical Properties: the characteristics that –The wayyg matter changes and behaves when determine how the composition of matter influenced by other matter and/or energy changes as a result of contact with other matter or the influence of energy – Characteristics that describe the behavior of matter

2 Chemical Properties Chemical Properties (cont.)

• One commonly cited chemical property is flammability, the ease with which a substance burns in a flame. Burning is a chemical reaction.

Classify Each of the following Classify Each of the following as a Physical or Chemical Property as a Physical or Chemical Property (cont.)

• Ethyl alcohol boiling at 78°C. • Ethyl alcohol boiling at 78°C. – Physical property: boiling point is a associated with a phase • Hardness of a diamond. change. It describes an inherent characteristic of alcohol. • Hardness of a diamond. • Sugar fermenting to form ethyl alcohol. – Physical property: describes an inherent characteristic of diamond – hardness

• Sugar fermenting to form ethyl alcohol. – Chemical property: describes behavior of sugar – forming a new substance (ethyl alcohol) through a chemical reaction

3 Changes in Matter Chemical Change

• Physical changes: changes to matter that do not result i n a ch ange th e f und ament al component s that make up the substance – State changes: boiling, melting, condensing

• Chemical changes: changes that involve a change in the fundamental components of the substance – Produce new substances – Chemical reactions occur – Reactants → Products

Classify Each of the following Chemical Change (cont.) as a Physical or Chemical Change

• Chemical change involves a chemical • Iron metal melting. reaction. At least one new substance is formed. • Iron combining with oxygen to form rust.

• Sugar fermenting to form ethyl alcohol.

4 Classify Each of the following Elements and Compounds as a Physical or Chemical Change (cont.)

• Iron metal melting. • Elements: substances that cannot be broken – Physical change: describes a state change, but the material is diilbdown into simpler substances b bhily chemical still iron reactions • Iron combining with oxygen to form rust. • Most substances are chemical combinations of – Chemical change: describes how iron and oxygen react to elements. These combinations are called make a new substance, rust compounds. – Compounds are made of elements • Sugar fermenting to form ethyl alcohol. – Compounds can be broken down into elements – Chemical change: describes how sugar forms a new – Properties of the compound not related to the properties of the substance (ethyl alcohol) via a chemical reaction elements that compose it – Same chemical composition at all times

Classification of Matter Pure Substances

Matter • Pure substances – All samples have the same physical and chemical Pure Substance Mixture properties Constant Composition Variable Composition Homogeneous – Constant composition: all samples have the same composition • Homogeneous: uniform throughout, appears to be one thing – Homogeneous – Pure substances – Solutions (homogeneous mixtures) – Separate into components based on chemical • Heterogeneous: non-uniform, contains regions with different properties properties than other regions

5 Mixtures Pure Substances vs. Mixtures

• Mixtures – Different samples may show different properties – Variable composition – Homogeneous or heterogeneous –Seppparate into components based on ppyhysical properties • All mixtures are made of pure substances

Solutions Solutions (cont.)

• A solution is a homogeneous mixture.

• Phase can be gaseous, liquid, or solid.

6 Identity Each of the following as a Pure Substance, Identity Each of the following as a Pure Substance, Homogeneous, Mixture, or Heterogeneous Mixture. Homogeneous Mixture, or Heterogeneous Mixture (cont.)

• Gasoline • Gasoline – A homogenous mixture • A stream with gravel on the bottom • A stream with gravel on the bottom • Copper metal – A heterogeneous mixture

• Copper metal – A pure substance (all elements are pure substances)

Separation of Mixtures Separation of Mixture (cont.)

• Mixtures can be separated based on different ppyhysical pro perties of the com ponents – Physical change Different Physical Property Technique Boiling point Distillation State of matter Filtra tion (solid/liquid/gas) Adherence to a surface Chromatography Volatility Evaporation

7 Separation of a Mixture (cont.) Another Look at Matter

Energy and Energy Changes Energy and Energy Changes

• Energy: ability to do work or produce heat • Potential Energy: – Chemical, mechanical, thermal, electrical , radiant , energy due to sound, nuclear composition or position – Potential and kinetic

• Energy may affect matter. • Kinetic Energy: energy – e.g. Raise its temperature, eventually causing a state due to motion changg,e, or cause a chemical chang e such as –-½mv½ mv2 decomposition

• All physical changes and chemical changes involve energy changes.

8 Energy and Energy Changes (cont.) Temperature and Heat

• Law of Conservation of Energy: energy can • Heat: a flow of energy due to a temperature be converted from one form to another, but difference cannot be created or destroyed • Temperature: a measure of the random motions of the components of a substance

Temperature and Heat (cont.) Exothermic vs. Endothermic

• System: that part of the universe that we wish to stdtudy • Surroundings: everything else in the universe • Exothermic process: a process that results in the evolution of heat - Example: when a match is struck, it is an exothermic ppgyprocess because energy is produced as heat. • Endothermic process: absorbs heat - Example: melting ice to form liquid water is an endothermic process because the ice absorbs heat in order to melt

9 Exothermic Process Units of Energy

• One calorie = amount of energy needed to raise the temperature of one gram of water by 1°C – kcal = energy needed to raise the temperature of 1000 g of water 1°C • joule – 4.184 J = 1 cal • In nutrition, calories are capitalized. – 1 Cal = 1 kcal

Example - Converting Calories to Joules Energy & Temperature of Matter

Convert 60.1 cal to • The amount the temperature of an object increases depends on the amount of heat added (q) . joules. – If you double the added heat energy the temperature will increase twice as much. 1 cal = 4.184 joules • The amount the temperature of an object increases 4.184 J when heat is added depends on its mass 60. 1cal × = 251J – If you dou ble t he mass it w ill tak e twi ce as much h eat 1 cal energy to raise the temperature the same amount.

10 Specific Heat Capacity Specific Heat Capacity

• Specific heat (s): the amount of energy required to raise the temperature of one gram of a substance by one degree Celsius

J By definition , the specific heat of water is 4.184 g °C

Amount of Heat = Specific Heat x Mass x Temperature Change Q = s x m x ∆T

Example #1: Example #1 (cont.)

J Calculate the amount of heat energy (in joules) Specific heat of water = 4.184 g °C needed to raise the temperature of 7.40 g of water from 29.0°C to 46.0°C. Mass = 7.40 g Temperature change = 46.0°C – 29.0°C = 17.0°C

Q = s • m • ∆T

J Heat = 4.184 ×7.40g×17.0°C = 526 J g °C

11 Example #2 Example #2

A 1.6 g sample of metal that appears to be Q = s× m× ∆T gold requi res 5 .8 J to rai se th e temperature Q s = from 23°C to 41°C. Is the metal pure gold? m × ∆T ∆T = 41°C - 23°C = 18°C 5.8 J J s = = 0.20 1.6 g x 18°C g °C J Table 10.1 lists the specific heat of gold as 0.13 g °C Therefore the metal cannot be pure gold.

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