Ch. 3: Matter and Energy

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

2. I. Chapter Outline • Introduction • Classifying Matter • Physical/Chemical Properties/Changes • Conservation of Matter • Energy • Temperature • Heat Capacity

3. I. Introduction • Everything around you is composed of matter. • Besides matter, energy is the other major component of our universe.

4. 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.

5. 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.

6. II. Atomic and Molecular Matter

7. II. Actual Images of Atoms and Molecules

8. 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

9. II. States of Matter

10. II. The Solid State

11. II. Properties of Different States

12. 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

13. II. Elements • element: a pure substance that cannot be broken down into simpler substances

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

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

16. II. Classification by Composition

17. II. Sample Problem • Classify the following as a pure substance or mixture. Further classify them as an element, compound, homogeneous, or heterogeneous. • blood • sugar • mercury in a thermometer • chicken noodle soup

18. 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

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

20. III. An Iron Nail Rusts • When iron rusts, it must react and incorporate oxygen to become a new compound.

21. III. Sample Problem • Identify the following as physical or chemical properties. • Hydrogen gas is explosive. • Silver has a shiny appearance. • Dry ice sublimes (goes from solid directly to vapor). • Copper turns green when exposed to air.

22. 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.

23. III. Physical/Chemical Changes

24. III. Sample Problem • Categorize the following as either a physical or chemical change. • Copper metal forming a blue solution when dropped in concentrated nitric acid. • A train flattening a penny. • A match igniting a firework. • Ice melting into liquid water.

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

26. 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

27. 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 object’s total energy is the sum of its PE and KE

28. 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.

29. 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.

30. V. Energy Unit Conversions • There are three common units for energy.

31. V. Sample Problem • The complete combustion of a wooden match produces about 512 cal of heat. How many kilojoules are produced?

32. 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.

33. V. Energy Diagrams • Chemical reactions can either be exothermic or endothermic. • exothermic: release energy to surroundings • endothermic: absorb energy from surroundings

34. V. Sample Problem • Identify the following changes as exothermic or endothermic. • Water freezing into ice. • Propane burning. • Isopropyl alcohol evaporating from skin.

35. 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?

36. 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.

37. VI. Temperature Scales

38. VI. Temperature Conversions • The formulas below allow conversion between different temperature units.

39. VI. Sample Problem • Convert 67 °F to kelvin and degrees Celsius.

40. 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

41. VII. Specific Heat Capacities

42. VII. Energy and Heat Capacity • Heat absorbed and temperature change are directly related as shown in the equation below.

43. 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.

44. 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.

45. 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.