How is matter classified?

Transcription

1 Matter How is matter classified? AP Chemistry (1 of 45) AP Chemistry (2 of 45) Solids Liquids AP Chemistry (3 of 45) AP Chemistry (4 of 45) Gases Classification Scheme for Matter AP Chemistry (5 of 45) AP Chemistry (6 of 45) Pure Substances Mixtures AP Chemistry (7 of 45) AP Chemistry (8 of 45) 1

2 How is matter classified? Matter Matter can be defined as anything that has mass and volume. Two principal ways of classifying matter are according to its physical state (as a gas, liquid, or solid) and according to its composition (as an element, compound, or mixture). Mass is a measure of the amount of matter in an object. Volume is the amount of three dimensional space an object takes up. Liquids Properties of Liquids No definite shape Definite volume The particles in a liquid are not as close together as the particles in a solid. The particles in a liquid slide over each other. Liquids pour easily. Solids Properties of Solids Definite shape Definite volume Particles are held tightly together, usually in definite arrangements. Particles vibrate in fixed positions Solids are rigid. There are strong attractive forces between the particles. Classification Scheme for Matter Gases Properties of Gases No definite shape No definite volume The particles are far apart and are moving at high speeds, colliding repeatedly with each other and the walls of the container. Gases can be compressed. Gases diffuse. The attractive forces between the particles are very weak. Mixtures Pure Substances A mixture is a physical blend of two or more pure substances, each of which retains its own identity and properties. Mixtures have variable compositions. Mixtures can be separated by physical means such as filtration, distillation, and chromatography. Mixtures can be heterogeneous or homogeneous. A pure substance has a fixed composition and differs from a mixture in that every sample of a given pure substance has exactly the same characteristic properties and composition. (Pure substances are homogeneous.) Elements and compounds are pure substances. 2

3 Elements Compounds AP Chemistry (9 of 45) AP Chemistry (10 of 45) Types of Elements Metals Types of Elements Nonmetals AP Chemistry (11 of 45) AP Chemistry (12 of 45) Types of Elements Metalloids Homogeneous Mixtures AP Chemistry (13 of 45) AP Chemistry (14 of 45) Heterogeneous Mixtures Physical Properties AP Chemistry (15 of 45) AP Chemistry (16 of 45) 3

4 Compounds Compounds are pure substances composed of two or more types of atoms. These atoms are always combined in the same ratio by mass. For example, pure water is always 11.2% hydrogen and 88.8% oxygen by mass. Compounds can be decomposed into simpler substances by chemical changes. For example, water can be decomposed by passing an electric current through it. This process is known as electrolysis. Examples: water, sodium chloride, sucrose Elements Elements are pure substances composed of only one type of atom. Elements cannot be decomposed into simpler substances by chemical or physical changes. Examples: gold, aluminum, oxygen, chlorine Types of Elements - Nonmetals Properties of Nonmetals 1. They are dull and brittle. 2. They don t conduct heat and electricity well. 3. They have relatively low boiling and freezing points. 4. They exist in all three phases at STP, but most are gases. The nonmetals are found on the right side of the periodic table (except for hydrogen). Examples: hydrogen, oxygen, nitrogen, carbon, sulfur, bromine Types of Elements Metals Properties of Metals 1. They are malleable, ductile and have luster. 2. They are good conductors of heat and electricity. 3. They have relatively high densities. 4. They are solids at standard temperature and pressure (STP), except for mercury, which is a liquid. 5. They have relatively high melting points, except for mercury and gallium. 6. They don t combine chemically with other metals. Metals combine physically to create alloys. Examples: lead, copper, sodium, lithium, aluminum Homogeneous Mixtures A homogeneous mixture has the same composition throughout. The components are indistinguishable. A homogeneous mixture is called a solution. Examples: air, sugar in water, stainless steel Types of Elements- Metalloids Properties of Metalloids 1. Metalloids have properties of both metals and nonmetals. 2. They are semiconductors at temperatures higher than room temperature. 3. They are all solids at STP. The metalloids are found between the metals and nonmetals (along the stair-step) on the periodic table Examples: Boron, Silicon, Germanium, Arsenic, Antimony Physical Properties The physical properties of a substance can be observed without any change in the composition of the substance. Examples: melting point, density, color Extensive physical properties depend on the amount of matter that is present. Examples: volume, mass Intensive physical properties do not depend on the amount of matter present. Examples: density, color, melting point Heterogeneous Mixtures A heterogeneous mixture does not have the same composition throughout. The components are distinguishable. Examples: Muddy river water, granite, wood, blood 4

5 Chemical Properties Physical Changes AP Chemistry (17 of 45) AP Chemistry (18 of 45) Changes in State Chemical Changes AP Chemistry (19 of 45) AP Chemistry (20 of 45) Indications of a Chemical Reaction Exothermic Reactions AP Chemistry (21 of 45) AP Chemistry (22 of 45) Endothermic Reactions Law of Conservation of Mass AP Chemistry (23 of 45) AP Chemistry (24 of 45) 5

6 Physical Changes Chemical Properties During physical changes a substance changes its physical appearance, but not its composition. Examples: folding paper, grinding, cutting, changes in state (melting, freezing, vaporization, condensation, deposition, sublimation) The chemical properties of a substance describe the way a substance may change or react to form other substances. Examples: flammability, reactivity with acid Chemical Changes A change in which one or more substances are converted into different substances is called a chemical change or chemical reaction. Examples: burning wood, food spoiling Changes in State Changes in state are examples of physical changes. The substances that react in a chemical change are called the reactants. The substances that are formed by the chemical change are called the products. 2Na + Cl 2 2NaCl reactants products Exothermic Reactions Reactions in which energy is released are called exothermic reactions. The energy that is released in an exothermic reaction was originally stored in the molecules of the reactants. CH 4 + 2O 2 CO 2 + 2H 2 O + energy Energy Content R P Reaction Progress Indications of a Chemical Reaction Certain easily observed changes usually indicate that a chemical reaction has occurred. 1. Energy is released or absorbed. 2. Production of a gas. 3. Formation of a precipitate. 4. Change in color or odor. Law of Conservation of Mass The Law of Conservation of Mass states that mass can be neither created nor destroyed by ordinary physical or chemical means. This means that the mass of the reactants must be equal to the mass of the products in a chemical reaction. Example: Hydrogen reacts with oxygen according to the following reaction to produce water. 2H 2 + O 2 2H 2 O How many grams of water will be produced if 4.0 g of hydrogen reacts with 32.0 g of oxygen? Ans g Endothermic Reactions Reactions in which energy is absorbed are called endothermic reactions. The energy that is absorbed is stored in the molecules of the products. 2NaCl + energy 2Na + Cl 2 Energy Content R P Reaction Progress 6

7 Qualitative vs. Quantitative Accuracy vs. Precision AP Chemistry (25 of 45) AP Chemistry (26 of 45) Error Accuracy, Precision and Error AP Chemistry (27 of 45) AP Chemistry (28 of 45) Calculating Percent Error Significant Figures AP Chemistry (29 of 45) AP Chemistry (30 of 45) Rounding Significant Figures Exact Numbers AP Chemistry (31 of 45) AP Chemistry (32 of 45) 7

8 Accuracy vs. Precision Qualitative vs. Quantitative Accuracy - refers to how close a measurement comes to the known or accepted value. Precision refers to the degree of agreement among several measurements of the same quantity. Precision reflects the reproducibility of a given type of measurement. Qualitative Measurements a measurement that gives descriptive, nonnumerical results. Ex. The solution is blue. Quantitative Measurements a measurement that gives results in a definite form, usually as numbers and units. Ex. The concentration of the solution is 0.5 M. Accuracy, Precision and Error Error Error refers to the deviation of a measurement from the known or accepted value low accuracy low accuracy high accuracy high precision low precision high precision small random error large random error small random error large systematic error no systematic error A random error (also called indeterminate error) means that an error has an equal probability of being high or low. This type of error occurs in estimating the value of the last digit of a measurement. Systematic error (or determinate error) occurs in the same direction each time; it is either always high or always low. Significant Figures Rules for Determining Significant Figures 1. Nonzero integers always count as significant figures. Ex. 125 has 3 significant figures 2. Leading zeros (zeros at the front of a number) do not count as significant figures. Ex has 2 significant figures. 3. Captive zeros (zeros between non zero digits) always count as significant figures. Ex has 4 significant figures 4. Trailing zeros (zeros at the right end of a number) may or may not be significant. They are significant only if the number contains a decimal point. Ex. 100 has 1 significant figure, 100. and 1.00 x 10 2 both have 3 significant figures. Percent Error = Calculating Percent Error Percent Error = Accepted Value- Experimental Value Accepted Value x100 =1.78% x100 Example: What is the percent error of a length measurement of cm if the accepted value is cm? Significant Figures Exact Numbers Many times calculations involve numbers that were not obtained using measuring devices but were determined by counting: 3 apples, 8 molecules. Such numbers are called exact numbers. Exact numbers can be assumed to have an infinite number of significant figures. Exact numbers do not limit the number of significant figures when used in a calculation. Rounding Rules for Rounding 1. In a series of calculations, carry the extra digits through to the final result, then round. 2. If the digit to be removed a. is less than 5, the preceding digit stays the same. For example, 1.33 rounded to two significant digits is 1.3. b. is equal to or greater than 5, the preceding digit is increased by 1. For example, 1.36 rounded to two significant digits is Remember, when rounding, use only the first number to the right of the last significant figure. Do not round sequentially. For example: a grade of is an 89 not a 90. 8

9 Significant Figures and Calculations Addition & Subtraction Significant Figures and Calculations Multiplication & Division AP Chemistry (33 of 45) AP Chemistry (34 of 45) Fundamental SI Units Derived SI Units AP Chemistry (35 of 45) AP Chemistry (36 of 45) Common SI Prefixes Mass vs. Weight AP Chemistry (37 of 45) AP Chemistry (38 of 45) Volume Density AP Chemistry (39 of 45) AP Chemistry (40 of 45) 9

10 Significant Figures and Calculations For multiplication or division, the number of significant figures in the result is the same as the number in the least precise measurement used in the calculation. For example consider the calculation significant figures x significant figures 6.38 = 6.4 (two significant figures) Significant Figures and Calculations For addition or subtraction, the result has the same number of decimal places as the least precise measurement used in the calculation. For example, consider the calculation decimal places decimal place decimal pl = 31.1 (one decimal place) Derived SI Units Derived units combinations of SI base units Fundamental SI Units Quantity area volume density molar mass Derivation length width length width height mass volume mass amount of substance Physical Quantity Name of Unit Abbreviation mass kilogram kg length meter m time second s temperature kelvin K amount of substance mole mol electric current ampere A luminous intensity candela cd Mass vs. Weight Common SI Prefixes Mass a measure of the amount of matter in an object. Mass is determined by comparing the mass of an object with a set of standard masses that are part of the balance. Mass is measured in the lab in grams (g). Weight a measure of the gravitational pull on matter. Weight is typically measured on a spring scale. Taking weight measurements involves reading the amount that an object pulls down on a spring. Weight is measured in Newtons (N). This SI system is based on powers of ten. The prefixes we will commonly use in this system are as follows: kilo k = thousand times the base unit deca da = times the base unit deci d = 0.1 1/10 the base unit centi c = /100 the base unit milli m = /1000 the base unit Density Density the ratio of mass to volume, or mass divided by volume. Density = mass volume Example: What is the density of an 84.7 g sample of an unknown substance if the sample occupies a volume of 49.6 cm 3? 84.7g Density = =1.71g / cm cm 3 Volume Volume the amount of space occupied by an object. The volume of a liquid can be measured using a graduated cylinder. The volume of a regularly shaped object can be measured using a metric ruler. Water displacement can be used to measure the volume of an irregularly shaped object. 10

11 Temperature Graphic Relationships Direct Relationship AP Chemistry (41 of 45) AP Chemistry (42 of 45) Graphic Relationships Inverse Relationship Conversion Factors AP Chemistry (43 of 45) AP Chemistry (44 of 45) Dimensional Analysis AP Chemistry (45 of 45) 11

12 Direct Relationship A direct relationship occurs where one variable increases as the other variable increases. Two quantities are directly proportional to each other if dividing one by the other gives a constant value. Temperature Temperature is a measure of the average kinetic energy of the particles in a sample of matter. The two temperature scales commonly used in chemistry are the Celsius scale and the Kelvin scale. Kelvin = C Conversion Factors A conversion factor is a ratio derived from the equality between two different units that can be used to convert from one unit to another. Example: The relationship between inches and centimeters can be used to write two conversion factors. 1inch 2.54 cm and 2.54 cm 1inch Inverse Relationship An inverse relationship occurs where one variable increases as the other variable decreases. Two quantities are inversely proportional to each other if their product is constant. The first conversion factor can be used to convert from cm to inches. The second conversion factor can be used to convert from inches to centimeters. Dimensional Analysis Dimensional analysis is used to convert from one unit to another. 1. To convert form one unit to another, use the equivalence statement that relates the two units. 2. Derive the appropriate conversion factor by looking at the direction of the required change (to cancel the unwanted units). 3. Multiply the quantity to be converted by the conversion factor to give the quantity with the desired units. Example: Calculate the number of seconds in 13.0 hours. 60 min 60sec hours = sec onds 1hr 1min 12