1 Chemistry: Methods and Measurement GENERAL CHEMISTRY General, Organic and Biochemistry 7th Edition Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

1.1 The Discovery Process Chemistry - The study of matter… Matter - Anything that has mass and occupies space A table A piece paper What about air? Yes, it is matter

1.1 The Discovery Process Chemistry: the study of matter its chemical and physical properties the chemical and physical changes it undergoes the energy changes that accompany those processes Energy - the ability to do work to accomplish some change 1.1 The Discovery Process

1.1 The Discovery Process THE SCIENTIFIC METHOD The scientific method - a systematic approach to the discovery of new information 2 Characteristics of the scientific process Observation Formulation of a question Pattern recognition Developing theories Experimentation Summarizing information 1.1 The Discovery Process

1.1 The Discovery Process

Three States of Matter 1.2 Matter and Properties gas - particles widely separated, no definite shape or volume 2. liquid - particles closer together, definite volume but no definite shape 3. solid - particles are very close together, define shape and definite volume 4 1.2 Matter and Properties

Three States of Water (a) Solid (b) Liquid (c) Gas (vapor)

Comparison of the Three Physical States 1.2 Matter and Properties

Physical property - is observed without changing the composition or identity of a substance Physical change - produces a recognizable difference in the appearance of a substance without causing any change in its composition or identity conversion from one physical state to another melting an ice cube 5 1.2 Matter and Properties

Separation by Physical Properties Magnetic iron is separated from other nonmagnetic substances, such as sand. This property is used as a large-scale process in the recycling industry.

Chemical property - result in a change in composition and can be observed only through a chemical reaction Chemical reaction (chemical change) - a process of rearranging, removing, replacing, or adding atoms to produce new substances 5 1.2 Matter and Properties hydrogen + oxygen  water reactants products

Classify the following as either a chemical or physical property: Color Flammability Hardness Odor Taste 1.2 Matter and Properties

Classify the following as either a chemical or physical change: Boiling water becomes steam Butter turns rancid Burning of wood Mountain snow pack melting in spring Decay of leaves in winter 1.2 Matter and Properties

Intensive properties - a property of matter that is independent of the quantity of the substance Density Specific gravity Extensive properties - a property of matter that is dependent on the quantity of the substance Mass Volume 6 1.2 Matter and Properties

Classification of Matter 1.2 Matter and Properties Pure substance - a substance that has only one component Mixture - a combination of two or more pure substances in which each substance retains its own identity, not undergoing a chemical reaction 7

Classification of Matter 1.2 Matter and Properties Element - a pure substance that cannot be changed into a simpler form of matter by any chemical reaction Compound - a substance resulting from the combination of two or more elements in a definite, reproducible way, in a fixed ratio 7

Classification of Matter 1.2 Matter and Properties Mixture - a combination of two or more pure substances in which each substance retains its own identity Homogeneous - uniform composition, particles well mixed, thoroughly intermingled Heterogeneous – nonuniform composition, random placement 7

Classes of Matter 1.2 Matter and Properties

1.3 Significant Figures and Scientific Notation Information-bearing digits in a number are significant figures The measuring device used determines the number of significant figures a measurement has The amount of uncertainty associated with a measurement is indicated by the number of digits or figures used to represent the information 8

1.3 Significant Figures and Scientific Notation Significant figures - all digits in a number representing data or results that are known with certainty plus one uncertain digit

Recognition of Significant Figures All nonzero digits are significant 7.314 has four significant digits The number of significant digits is independent of the position of the decimal point 73.14 also has four significant digits Zeros located between nonzero digits are significant 60.052 has five significant digits 1.3 Significant Figures and Scientific Notation

Use of Zeros in Significant Figures Zeros at the end of a number (trailing zeros) are significant if the number contains a decimal point. 4.70 has three significant digits Trailing zeros are insignificant if the number does not contain a decimal point. 100 has one significant digit; 100. has three Zeros to the left of the first nonzero integer are not significant. 0.0032 has two significant digits 1.3 Significant Figures and Scientific Notation

How many significant figures are in the following? 3.400 3004 300. 0.003040 1.3 Significant Figures and Scientific Notation

1.3 Significant Figures and Scientific Notation Used to express very large or very small numbers easily and with the correct number of significant figures Represents a number as a power of ten Example: 4,300 = 4.3  1,000 = 4.3  103 1.3 Significant Figures and Scientific Notation

1.3 Significant Figures and Scientific Notation For numbers greater than 1, the original decimal point is moved x places to the left, and the resulting number is multiplied by 10x The exponent x is a positive number equal to the number of places the decimal point moved 5340 = 5.34  103 What if you want to show the above number has four significant figures? = 5.340  103 1.3 Significant Figures and Scientific Notation

1.3 Significant Figures and Scientific Notation To convert a number less than 1 to scientific notation, the original decimal point is moved x places to the right, and the resulting number is multiplied by 10–x The exponent x is a negative number equal to the number of places the decimal point moved 0.0534 = 5.34  10–2 1.3 Significant Figures and Scientific Notation

Convert the following to 2 SF using scientific notation… 0.0024 48.20 224 0.0180

1.3 Significant Figures and Scientific Notation Types of Uncertainty Error - the difference between the true value and our estimation Random Systematic Accuracy - the degree of agreement between the true value and the measured value Precision - a measure of the agreement of replicate measurements 1.3 Significant Figures and Scientific Notation

Exact and Inexact Numbers Inexact numbers have uncertainty by definition (any measured value) Exact numbers are a consequence of counting A set of counted items (beakers on a shelf) has no uncertainty Exact numbers by definition have an infinite number of significant figures 1.3 Significant Figures and Scientific Notation

Rules for Rounding Off Numbers When the number to be dropped is less than 5 the preceding number is not changed When the number to be dropped is 5 or larger, the preceding number is increased by one unit Round the following number to 3 significant figures: 3.34966  104 1.3 Significant Figures and Scientific Notation =3.35  104

Round off each number to three significant figures: 61.40 6.171 0.066494 1.3 Significant Figures and Scientific Notation

1.4 Units and Unit Conversion Units - the basic quantity of mass, volume or whatever is being measured A measurement is useless without units 9 English system - a collection of functionally unrelated units Difficult to convert from one unit to another 1 foot = 12 inches = 0.33 yard = 1/5280 miles Metric System - composed of a set of units that are related to each other decimally, systematic Units relate by powers of ten

1.4 Units and Unit Conversion Basic Units of the Metric System Mass gram g Length meter m Volume liter L 1.4 Units and Unit Conversion Basic units are the units of a quantity without any metric prefix.

1.4 Units and Unit Conversion

1.4 Units and Unit Conversion You must be able to convert between units within the metric system between the English system and metric system The method used for conversion is called the Factor-Label Method (FLM) or Dimensional Analysis 9 1.4 Units and Unit Conversion !!!!!!!!!!! VERY IMPORTANT !!!!!!!!!!!

1.4 Units and Unit Conversion Let your units do the work for you by simply memorizing connections between units. For example: How many donuts are in one dozen? 1.4 Units and Unit Conversion We say: “Twelve donuts are in a dozen.” Or: 12 donuts = 1 dozen donuts What does any number divided by itself equal? ONE!

1.4 Units and Unit Conversion This fraction is called a unit or conversion factor 1.4 Units and Unit Conversion Multiplication by a unit factor does not change the amount – only the unit

1.4 Units and Unit Conversion Example: How many donuts are in 3.5 dozen? You can probably do this in your head but try it using the Factor-Label Method: 1.4 Units and Unit Conversion

1.4 Units and Unit Conversion Start with the given information... 3.5 dozen = 42 donuts 1.4 Units and Unit Conversion Then set up your unit factor... See that the units cancel... Then multiply and divide all numbers...

Common English System Units 1.4 Units and Unit Conversion Convert 12 gallons to quarts Convert 175 lbs to ounces Convert 10,000 ft to miles

Intersystem Conversion Units 1.4 Units and Unit Conversion Convert 4.00 ounces to kilograms Convert 175 lbs to kgs Convert 100 yards to meters

Practice Unit Conversions Convert 12 gallons to quarts Convert 10 cm to meters Convert 360 feet to miles Convert 60 miles/hr to m/s 5. Convert 1.8 in2 to cm2 1.4 Units and Unit Conversion

The speed of light is 299,792,458 m/s  convert this to miles per hour. The speed of sound is 340.29 m/s  convert this to miles per hour (Use: 1 mile = 1.6 km; 1 km = 1000 m; 60 s = 1 min; 60 min = 1 hour)

1.5 Experimental Quantities Mass - the quantity of matter in an object not synonymous with weight standard unit is the gram Weight = mass  acceleration due to gravity Mass must be measured on a balance (not a scale)

1.5 Experimental Quantities Units should be chosen to best suit the quantity described: A dump truck is measured in tons A person is measured in kg or pounds A paperclip is measured in g or ounces For atoms, we use the atomic mass unit (amu) 1 amu = 1.661  10-24 g 1.5 Experimental Quantities

1.5 Experimental Quantities Length - the distance between two points standard unit is the meter long distances are measured in km distances between atoms are measured in nm, 1 nm = 10-9 m Volume - the space occupied by an object standard unit is the liter the liter is the volume occupied by 1000 grams of water at 4 oC 1 mL = 1/1000 L = 1 cm3 1.5 Experimental Quantities

1.5 Experimental Quantities The milliliter (mL) and the cubic centimeter (cm3) are equivalent 1.5 Experimental Quantities

1.5 Experimental Quantities Time metric unit is the second Temperature - the degree of “hotness” of an object 10 1.5 Experimental Quantities

Conversions Between Fahrenheit and Celsius 1.5 Experimental Quantities Convert 75oC to oF Convert -10oF to oC

Kelvin Temperature Scale The Kelvin scale is another temperature scale. It is of particular importance because it is directly related to molecular motion. As molecular speed increases, the Kelvin temperature proportionately increases. 1.5 Experimental Quantities K = oC + 273 Convert 20ºC to K

Density and Specific Gravity the ratio of mass to volume an intensive property use to characterize a substance as each substance has a unique density Units for density include: g/mL g/cm3 g/cc 11 1.5 Experimental Quantities

1.5 Experimental Quantities cork 1.5 Experimental Quantities water brass nut liquid mercury

Calculating the Density of a Solid 2.00 cm3 of aluminum are found to weigh 5.40g. Calculate the density of aluminum in units of g/cm3. 1.5 Experimental Quantities

1.5 Experimental Quantities Density Calculations Air has a density of 0.0013 g/mL. What is the mass of 6.0-L sample of air? Calculate the mass in grams of 10.0 mL if mercury (Hg) if the density of Hg is 13.6 g/mL. Calculate the volume in milliliters, of a liquid that has a density of 1.20 g/mL and a mass of 5.00 grams. 1.5 Experimental Quantities

1.5 Experimental Quantities Specific Gravity Values of density are often related to a standard Specific gravity - the ratio of the density of the object in question to the density of pure water at 4oC Specific gravity is a unitless term because the 2 units cancel Often the health industry uses specific gravity to test urine and blood samples 1.5 Experimental Quantities