AP Chemistry Chapter 1

1. AP ChemistryChapter 1 Chemistry: The Study of Change

2. About Chemistry • Chemistry is the study of matter and the changes it undergoes. • Matter is defined as anything that has mass and occupies space. • The three physical states of matter are solid, liquid, and gas. • All matter exists in one or another of these three states, depending on the temperature and pressure of the surrounding environment.

5. A solid has definite volume and definite shape. A liquid has a definite volume but not a definite shape. A gas does not have a definite volume or shape.

6. About Chemistry • Chemists are concerned with developing the tools used to study matter and the given forms of matter into new and different substances, and to the discovery of the properties and uses of these new materials. • Chemists usually observe matter and the changes it undergoes in the macroscopic world. This refers to the objects we can see and touch, and deal with everyday. • However, our interpretations of matter involve atoms and molecules and their properties. Because atoms and molecules are so extremely small, we refer to them as belonging to the microscopic world.

7. Matter and Its Properties • A pure substance is a form of matter that has definitecomposition and distinct properties. • Examples are water, table salt, and iron. • Just as each individual person has a set of characteristics, such as fingerprints and color of eyes and hair, each pure substance has characteristic properties.

8. Pure Substances • There are two types of pure substance: elements and compounds. • An element is a pure substance that cannot be decomposed into simpler substances by ordinarychemicalreactions. • Elements are the building blocks of which all compounds are composed. • Nitrogen, oxygen, and iron are examples of elements. • Compounds are pure substances that are composed of two or moreelements combined in definiteproportions. • Compounds can be brokendown into the elements of which they are composed by chemicalmeans. • Water and table salt are compounds.

11. Pure Substances • The action of an electric current, called electrolysis, is one method that can be used to decompose both water and molten table salt into their constituent elements. • Pure water consists of 89 % oxygen and 11 % hydrogen by mass. • Pure salt contains 39% sodium and 61% chlorine.

13. Mixtures • Pure substances can be brought together to form mixtures. • Mixtures are combinations of two more substances with variablecomposition. • They can be homogeneous or heterogeneous depending on the state of subdivision of the components.

14. Homogeneous Mixtures • Salt water is a uniform mixture of table salt (NaCl) and water. • The original crystals of salt have dissolved and are dispersed evenly. • On the ordinary scale of observation we cannot detect any chemical or physical differences between adjacent regions of the mixture. • The particles of salt are too small to observe. • A mixture that has the same composition throughout is said to be a homogeneous mixture.

15. Homogeneous Mixtures • The properties of a homogeneous mixture vary since they depend on the percent composition. • For example, the hardness of steel, a solid mixture of iron and carbon, depends on the percentage of carbon that is added to iron. • Homogeneous mixtures are also called solutions.

16. Heterogeneous Mixtures • Heterogeneous mixtures are notuniform in composition. • And indeed the individual particles of their components can often be seen by the unaidedeye. • For example, when preparing home-made ice cream, you use a mixture of ice and rock salt. • This is a heterogeneous mixture. • The individual chunks (particles) of ice and salt are clearly visible, and the particles are so large, they are not evenly dispersed.

17. Heterogeneous Mixtures • The composition of this mixture varies from place to place within the mixture itself. • Naturalair is actually a heterogeneous mixture. • It consists of nitrogen, oxygen, and argon gases, but also contains solid particles of pollen and dust. • Any mixture, whether it be homogeneous or heterogeneous, can be separated into its pure components by physicalmeans.

20. Physical and Chemical Properties of Matter • Physical properties are those properties that can be measured and observed withoutchanging the identity or composition of the substance. • Physical properties include color, hardness, solubility, density, specific heat, melting point, and boiling point. • Physical changes are those that take place with no change in chemical composition. • Changes of a substance from one state of matter to another do not change its chemical composition and are examples of physical changes. • The three forms of water we call ice, liquid water, and steam are all the same substance, just different physical states.

21. Physical and Chemical Properties of Matter • Chemical properties most often are descriptions of reactions that a substance undergoes when brought in contact with other substances. • In a chemical reaction the original substance or substances are changed into newsubstances. • When sodium metal and chlorine gas are heated together, a white solid called sodium chloride is formed. • That this is a chemical change is evident by the observation that sodium, a shiny metal, and chlorine, a pale yellow green gas, have disappeared, and in their place is a substance with a completely new set of properties. • Sodium chloride is a white solid that melts at very high temperatures.

25. Physical and Chemical Properties of Matter • All properties of matter are either extensive or intensive properties. • Extensive properties depend on the amount of matter being considered. • Volume and mass are examples. • In contrast, temperature and density are two properties that do not depend on the amount of mass present. • Thus, they are intensive properties.

27. Units of Measurement • Since 1960 a coherent system of units, known as the SI, has been in effect and is accepted among scientists and engineers. • SI is the abbreviation for Le Systeme International d’Unites. • The SI has as its base units: the kilograms (kg) for mass, the meter (m) for length, the second (s) for time, the Kelvin(K) for temperature, and the mole (mol) for amount of substance.

29. Units of Measurement • Combinations of base units produce derivedunits. For example, velocity is distance traveled per unit of time. • Therefore, velocity has units of meters per second (m/s) and is a derived unit rather than a base unit. • A major advantage of the SI is that it uses the decimal system. • A list of the SI prefixes is given in Table 1.3. • These prefixes will be used throughout your study of chemistry. • Note that in the SI the same prefix can be applied to any base unit or derived unit.

30. Units of Measurement • Thus the prefix milli can be used to describe a unit that is 1/1000 of a gram, or 1/1000 of a meter, or 1/1000 of any SI unit. 1 milligram = 1/1000 gram1 millimeter = 1/1000 meter1 millisecond = 1/1000 second

32. Volume • Volume is a derived unit. It can be expressed in terms of lengthcubed because for rectangular solids volume is equal to length times width times height all of which have the base unit meters. volume = length x width x height = length3volume = m x m x m = m3 • The SI unit for volume is the cubic meter (m3) which is the volume of a cube 1 m on each edge. • This unit is too large and related units such as the cubic centimeter (cm3) and cubic decimeter (dm3) are often used.

34. Volume • The volume of a fluid is usually measured in liters (L). • A liter is roughly the size of a quart (1.06 qt). A liter is the volume of a cube with an edge of 10 cm, and therefore is 1000 cm3. • The volume unit you are most likely to use in chemistry is the milliliter (mL). One mL is 0.001 L, and so 1000 mL equal 1 L. • Since 1 L = 1000 cm3 = 1000 mL, then one milliliter (mL) is equal to one cubic centimeter (cm3). 1 mL = 1 cm3

37. Density • Density is an important physical property of objects and substances. • The density of an object is defined as the ratio of its mass to its volume. • Density has units that are derived from the base units for mass and volume which are kilograms per cubic meter. • However, it is more convenient to report densities in units of grams per cubic centimeter or grams per milliliter.

39. Temperature • Chemists use two temperature scales, the Kelvin scale (K) and the Celsius scale (°C). • A third scale, the Fahrenheit scale (°F), is commonly used in the United States. • The Celsius scale defines 0°C as the freezingpoint of water and 100°C as the boilingpoint of water.

41. Handling Numbers • Many of the quantities (numbers) encountered in chemistry are more easily manipulated when they are written in a form known as scientific or exponentialnotation. • Very large or very small numbers are expressed as N x 10n, where N is a number between 1 and 10 and n corresponds to the exponents 1, 2, 3, etc.

42. Handling Numbers • Recall that the exponent tells you how many times to multiply the number 10 by itself. Therefore: 103 = 10 x 10 x 10= 1000 102 = 10 x 10= 100 101 = 10= 10 100= 1

43. Handling Numbers • To write the number 5000 in exponential notation, we could rewrite it as 5 x 1000. But 1000 can be written 10 x 10 x 10, or just 103. Therefore, 5000 can be written as 5 x 103. • More simply, to find the exponent of 10, just count the number of places that the decimal point must be moved to the left to give the number N. • For the number 5000, moving the decimal point three places to the left gives 5 x 103.

44. Handling Numbers • Fractional numbers can also be expressed in scientific notation, but in this case the decimal point will be moved to the right and the exponent will be a negative number (–n). • For the number 0.05, the decimal point must be moved two places to the right to give N (a number between 1 and 10). • In exponential notation 0.05 becomes 5 x 10–2.

45. Handling Numbers • Multiplication and division of numbers that are expressed in exponential notation are accomplished by operating on the coefficients (N) and exponentials (n) separately. • Recall when multiplying two exponential numbers that the exponents are added. • When dividing exponential numbers, we subtract the exponent in the denominator from the exponent in the numerator.

46. Significant Figures • Measurements, and calculations based on measurements, must be reported so as to convey information about the number of meaningful digits. • The significant figures are those digits in a measured number that include all certain digits plus one having uncertainty.

47. Significant Figures • If we measure the length of a desktop with a meterstick calibrated in millimeters and find the length to be 972.5 mm, then the digits 9, 7, and 2 mean there are 9 hundreds, 7 tens, and 2 ones, and because of the calibrations, these digits are certain. • Five-tenths, however, is an estimate because it falls between the finest calibrations. • Someone else might read it as 0.6 or 0.4. • The result could be expressed with error limits as 972.5 ± 0.1 mm. • This measurement provides three certain digits and one uncertain digit. • All of these provide useful information, and we say the number 972.5 has four significant figures.

49. Significant Figures • Always take care that the numbers you write reflect the proper number of meaningful digits. • The process of determining the correct number of significant figures after a calculation depends on the type of calculation. • First let us review the rules for expressing significant figures.

50. Guidelines for Writing Significant Figures • To determine the number of significant digits that are present in a written number, use the following rules: • Count all numbers significant except: • Leading zeros 0.0002 • Trailing zeros with a decimal point 2500