Ch. 3: Matter and Energy - San Diego Miramar College

Ch. 3: Matter and Energy - San Diego Miramar College

Ch. 3: Matter and Energy Dr. Namphol Sinkaset Chem 152: Introduction to General Chemistry I. Chapter Outline I.


Introduction Classifying Matter Physical/Chemical Properties/Changes Conservation of Matter Energy Temperature Heat Capacity

I. Introduction Everything around you is composed of matter. Besides matter, energy is the other

major component of our universe. II. Matter Matter is anything that occupies space and has mass. Some matter is easy to see (water,

wood), others are difficult (air, dust). The most basic building block of matter is the atom. II. Atoms and Molecules atoms: submicroscopic particles that are the fundamental building blocks of

all matter. Sometimes, atoms are bonded together to form molecules. molecules: two or more atoms joined to one another in specific geometric arrangements.

II. Atomic and Molecular Matter II. Actual Images of Atoms and Molecules II. States of Matter

Matter can be classified by its state. solid: closely-packed particles with fixed locations liquid: closely-packed particles, but free to move around gas: great distances between particles with free movement

II. States of Matter II. The Solid State II. Properties of Different States

II. Pure Substances and Mixtures Matter can be classified by its composition. pure substance: matter composed of only one type of atom or molecule mixture: matter composed of two or

more different types of atoms or molecules which may vary in proportion II. Elements element: a pure substance that cannot be broken down into simpler substances

II. Compounds compound: a pure substance composed of two or more elements in fixed definite proportions.

II. Mixtures Most matter exists in this form. heterogeneous: varied composition from one region to another homogeneous: uniform composition throughout II. Classification by

Composition II. Sample Problem Classify the following as a pure substance or mixture. Further classify them as an element, compound, homogeneous, or heterogeneous.

a) b) c) d) blood sugar

mercury in a thermometer chicken noodle soup III. Distinguishing Matter We use physical and chemical properties to tell the difference between samples of matter.

physical property: a property a substance displays without changing its composition chemical property: a property a substance displays only by changing its composition

III. Boiling Point of Water At the boiling point, water is converted to steam, but steam is just a different form of water.

III. An Iron Nail Rusts When iron rusts, it must react and incorporate oxygen to become a new compound. III. Sample Problem Identify the following as physical or chemical properties.

a) Hydrogen gas is explosive. b) Silver has a shiny appearance. c) Dry ice sublimes (goes from solid directly to vapor). d) Copper turns green when exposed to air. III. Physical/Chemical Changes

Physical/chemical changes are closely related to definitions of physical/chemical properties. physical change: matter changes its appearance, but not its composition chemical change: matter changes its composition

Chemical changes occur through chemical reactions in which reactants become products. III. Physical/Chemical Changes III. Sample Problem

Categorize the following as either a physical or chemical change. a) Copper metal forming a blue solution when dropped in concentrated nitric acid. b) A train flattening a penny. c) A match igniting a firework. d) Ice melting into liquid water.

IV. There is No New Matter In ordinary chemical reactions, matter is neither created nor destroyed. Known as Conservation of Mass. V. Energy

Physical and chemical changes are accompanied by energy changes. energy: the capacity to do work work: results from a force acting on a distance V. Two Types of Energy potential energy (PE): energy due to the

position or composition of the object kinetic energy (KE): energy due to motion of the object An objects total energy is the sum of its PE and KE V. Energy Conversions

The Law of Conservation of Energy states that energy is neither created nor destroyed. Energy can change from one form to

another or transferred from one object to another. V. Specific Types of Energy Electrical energy is the energy associated with the flow of electrical

charge. Thermal energy is the energy associated with motions of particles of matter. Chemical energy is a form of PE associated with positions of particles in a chemical system.

V. Energy Unit Conversions There are three common units for energy. V. Sample Problem The complete combustion of a wooden match produces about 512 cal of heat.

How many kilojoules are produced? V. System and Surroundings When describing energy changes, we need reference points. system: object of study surroundings: everything else

Systems with high PE tend to change such that their PE is lowered. V. Energy Diagrams Chemical reactions can either be exothermic or endothermic.

exothermic: release energy to surroundings endothermic: absorb energy from surroundings V. Sample Problem

Identify the following changes as exothermic or endothermic. a) Water freezing into ice. b) Propane burning. c) Isopropyl alcohol evaporating from skin. VI. Thermal Energy

Atoms and molecules of matter are in constant, random motion, which is the source of thermal energy. More motion = more thermal energy. Is there a way to easily measure this motion?

VI. Temperature and Heat Temperature is the measure of the thermal energy of a substance. The hotter an object, the greater the motion of its particles, and the greater the thermal energy. Heat is the transfer or exchange of

thermal energy caused by a temperature difference. VI. Temperature Scales VI. Temperature Conversions The formulas below allow conversion between

different temperature units. VI. Sample Problem Convert 67 F to kelvin and degrees Celsius. VII. Heating a Substance

When you heat a substance, its temperature changes. The amount of change depends on the substance. heat capacity: quantity of heat needed to raise the temp of substance by 1 C specific heat capacity: quantity of heat

needed to raise temp of 1 g of substance by 1 C VII. Specific Heat Capacities VII. Energy and Heat Capacity Heat absorbed and temperature change are

directly related as shown in the equation below. VII. Sample Problem Calculate the heat necessary to warm a 3.10 g sample of copper from -5.0 C to 37.0 C if the specific heat capacity of

copper is 0.385 J/g C. VII. Sample Problem A sample of lead (C = 0.128 J/g C) absorbs 11.3 J of heat, rising in temperature from 26 C to 38 C. Find the mass of the sample in grams.

VII. Sample Problem A 328-g sample of water absorbs 5.78 kJ of heat. If the water sample has an initial temperature of 35.3 C, what will be its final temperature? Note that C = 4.18 J/g C for water.

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