Chapter 1

1. Chapter 1 The Science of Chemistry Vocabulary Headings Important Information

2. Section 1: What is Chemistry? • Objectives: • Describe ways in which chemistry is a part of your daily life. • Describe the characteristics of three common states of matter. • Describe physical and chemical changes, and give examples of each. • Identify the reactants and products in a chemical reaction. • List four observations that suggest a chemical change has occurred.

3. Working with the Properties and Changes of Matter • Chemistry plays a vital role in your daily life and in the complex workings of your world. • Chemistry is concerned with the properties of chemicals and with the changes chemicals can undergo. • A chemical is any substance that has a definite composition. • Some chemicals exist naturally, others are manufactured.

4. You Depend on Chemicals Every Day • Why do many people think of chemicals in a negative way? • Mostly chemicals are discussed when we are talking about negative situations, for example: • Causes of pollution • Causes of explosions • Treatments dealing with cancer

5. Chemical Reactions Happen All Around You • Chemical reactionsare changes in chemicals. • They take place all around us and inside us. • Chemical reactions are necessary for living things to grow and for dead things to decay. • What are some examples of chemical reactions? • Cooking food • Taking a photo • Striking a match • Switching on a flash light • Starting a gasoline engine

6. Physical States of Matter • All matter is made of particles. • The type and arrangement of the particles in a sample of matter determine the properties of the matter. • Most of the matter you encounter is in one of three states of matter: solid, liquid, or gas. • Macroscopic refers to what you see with the unaided eye. • Microscopic refers to what you would see if you could see individual atoms

7. Properties of the Physical States • Solids have fixed volume and shape because their particles are held tightly in a rigid structure. They vibrate only slightly. • Liquids have fixed volume but not a fixed shape because their particles slip past one another. A liquid can flow and take the shape of its container. • Gases have neither fixed volume nor fixed shape. Gases will fill any container they occupy as their particles move apart. • There are other states of matter, that we will not discuss.

8. Changes of Matter • Many changes of matter happen: • Ice cube melts • Bicycle spokes rust • A red shirt fades • Water fogs a mirror • Milk sours • Two broad categories are: Physical changes & Chemical changes

9. Physical Changes • Physical changes are changes in which the identity of a substance doesn’t change. However, the arrangement, location, and speed of the particles that make up the substance may change. • Changes of state are physical changes. • Examples are: • Sugar dissolves in tea • Crushing a rock

10. Chemical Changes • In a chemical change, the identities of substances change and new substances form. Mercury(ii) oxide mercury + oxygen • In an equation, the substances on the left-hand side of the arrow are the reactants. They are used up in the reaction. • Substances on the right side of the arrow are the products. They are made by the reaction. • A chemical reaction is a rearrangement of the atoms that make up the reactant or the reactants. After rearrangement, those same atoms are present in the product or products. Atoms are not destroyed or created, so mass does not change during a chemical reaction.

11. Evidence of Chemical Change • Evidence that a chemical change may be happening: • The Evolution of a Gas: The production of a gas is often observed by bubbling, or by a change in color. • The Formation of a Precipitate: When two clear solutions are mixed and become cloudy, a precipitate has formed • The Release or Absorption of Energy: A change in temperature or the giving off of light energy are signs of an energy transfer. • A Color Change in the Reaction System: Look for a different color when two chemicals react. • The more of these signs you observe, the more likely a chemical change is taking place. • Some physical changes also have one or more of these signs so BE CAREFUL!!!

12. Section 2: Describing Matter • Objectives • Distinguish between different characteristics of matter, including mass, volume, and weight. • Identify and use SI units in measurements and calculations. • Set up conversion factors, and use them in calculations. • Identify and describe physical properties, including density. • Identify chemical properties.

13. Matter has Mass and Volume • Matteris anything that has mass and volume.

14. The Space an Object Occupies Is Its Volume • Volume is the space an object occupies. • The method used to determine volume depends on the nature of the matter being examined. • Sometimes you can measure volume by multiplying length, width, and height. • Graduated cylinders are used to measure the volume of a liquid. • The volume of a gas is the same as that of the container it fills.

15. The Quantity of Matter is the Mass • The mass of an object is the quantity of matter contained in that object. • Devices used for measuring mass in a laboratory are called balances.

16. Mass is Not Weight • Mass is related to weight but the two are not identical. • As long as the subject is not changed, it will have the same mass, no matter where it is in the universe. • The weight of that object is affected by its location in the universe because weight depends on gravity. • Weight is defined as the force produced by gravity acting on mass • The force that gravity exerts on an object is proportional to the object’s mass

17. Units of Measurement • In chemistry we prefer to describe properties of substances in quantitative terms, that is with numbers.

18. SI Base Units

19. Identify the SI base units and compare the base units to derived units • Derived units are made up of more than one base unit • Examples of derived units : g/cm3, g/mol, m3 • If it is not on the list of base units, it is a derived unit.

20. Define the main prefixes used in the metric system.

21. Define the main prefixes used in the metric system.

22. Converting One Unit to Another • In chemistry, you often need to convert a measurement from one unit to another. One way of doing this is to use a conversion factor. • A conversion factor is a simple ratio that relates two units that express a measurement of the same quantity. • An example of this is on page 13 of your book.

23. Convert between different metric units. • G _ _ M _ _ k h da (base) d c m _ _ m _ _ n • Start to the right of the given • Move the decimal to the right of the unknown • Add zeroes to hold the decimal place • 2km = ____________ cm • Move from the right of the k to the right of the c • 5 spaces to the right, need 5 zeroes to hold the place • 2km = 200,000cm

24. Convert the following measurements • 6.7 nm = ________m • 37 kg = ________cg • 23 ml = ________ kl • 9.7 dam = ______m • 0.0054 cg = _____mg • 0.5 cm = _______mm • 0.68 Mg = _______cg • 1Gm= _________nm • 300 mm = _______m • 2 km = ________hm • 0.90 cm = ______mm • 5.67 dm = ______km • 3.6 Gm = ________m • 4.5 mg= _________g • 34 kg = ________mg • 45 ms = ________s

25. Properties of Matter • The properties of a substance may be classified as physical or chemical.

26. Physical Properties • A physical property is a property that can be determined without changing the nature of the substance. • Color and state are example of physical properties. • Melting point and boiling point are also examples of physical properties.

27. Solve density problems. • Density is the Mass divided by the Volume • D = M/V • The equation can be rearranged to solve for any of the three variables. • M = D x V • V = M/D • Example • A block of aluminum occupies a volume of 15.0 mL and weighs 40.5 g. What is its density? • M = 40.5 g • V = 15.0 mL • D = M/V D = 40.5g/15.0 mL D = 2.7 g/mL

28. Density Problems cont. • What is the weight of the ethyl alcohol that exactly fills a 200.0 mL container? The density of ethyl alcohol is 0.789 g/mL. • D = 0.789 g/ml (3 sig figs) • V = 200.0 mL (4 sig figs) • M = D x V M = 0.789g/ml x 200.0 mL • M = 158 g (3 sig figs)

29. Density Problems cont. • What volume of silver metal will weigh exactly 2500.0 g. The density of silver is 10.5 g/cm3. • D = 10.5 g/cm3 • M = 2500.0 g • V = M / D V = 2500.0 g/ 10.5 g/cm3 • V = 238 cm3 (3 sig figs)

30. Density Problems cont. • A rectangular block of copper metal weighs 1896 g. The dimensions of the block are 8.4 cm by 5.5 cm by 4.6 cm. From this data, what is the density of copper? • First calculate the volume. (V = l x w x h) • V = 8.4 cm x 5.5 cm x 4.6 cm V = 212.52 cm3 • There are only 2 sig figs in the problem so you must round to 210 cm3 • Now calculate the density. • M = 1896 g • V = 210 cm3 • D = M/V D = 1896 g/ 210 cm3 D = 9.03 g/cm3 • Can only have 2 sig figs therefore the answer would be 9.0 g/cm3

31. Density Practice • Mercury metal is poured into a graduated cylinder that holds exactly 22.5 mL. The mercury used to fill the cylinder weighs 306.0 g. From this information, calculate the density of mercury. • A flask that weighs 345.8 g is filled with 225 mL of carbon tetrachloride. The weight of the flask and carbon tetrachloride is found to be 703.55 g. From this information, calculate the density of carbon tetrachloride. • Calculate the density of sulfuric acid if 35.4 mL of the acid weighs 65.14 g. • Find the mass of 250.0 mL of benzene. The density of benzene is 0.8765 g/mL. • A block of lead has dimensions of 4.50 cm by 5.20 cm by 6.00 cm. The block weighs 1587 g. From this information, calculate the density of lead. • What is the volume of a substance with a mass of 0.35 g and a density of 0.9 g/ml?

32. Density is the Ratio of Mass to Volume • The mass and volume of a sample are physical properties that can be determined without changing the substance. • The density of an object is another physical property: the mass of that object divided by its volume. • Density=mass/volume or D=m/v

33. Density Can be Used to Identify Substances • Because the density of a substance is the same for all samples, you can use this property to help identify substances. • One way to find the volume is to use the technique of water displacement.

34. Chemical Properties • What happens when matter has the chance to react with other kinds of matter are called chemical properties. • Chemical properties can only be identified by trying to cause a chemical change. • Examples: • Reactivity • Combustibility • Flammability

35. Section 3: How is Matter Classified • Objectives: • Distinguish between elements and compounds. • Distinguish between pure substances and mixtures. • Classify mixtures as homogeneous or heterogeneous. • Explain the difference between mixtures and compounds.

36. Classifying Matter • Everything is made of matter. • Atom: smallest unit of an element that maintains the properties of that element • All matter is composed of about 110 different kinds of atoms. • Even the biggest atoms are so small that it would take more than 3 million of them side by side to span just 1 millimeter.

37. Benefits of Classification • Because matter exists in so many different forms, having a way to classify matter is important for studying it.

38. Pure Substances • A pure substanceis a sample of matter, either a single element or a single compound, that has definite chemical and physical properties. • There are two types of pure substances: elements and compounds

39. Elements are Pure Substances • Elements are pure substances that contain only one kind of atom. • Each element has its own unique set of physical and chemical properties. • Each element is represented by a distinct chemical symbol.

40. Elements as Single Atoms or as Molecules • Some elements exist as single atoms. • An example of this is the helium gas in a balloon. • Because the gas exists in individual atoms, helium gas is known as a monatomic gas. • Other elements exist as molecules. • A molecule usually consists of two or more atoms combined in a definite ratio. • If an element consists of molecules, those molecules contain just one type of atom. • Oxygen and nitrogen are examples of diatomic elements because they exist as 2 atoms joined together.

41. Some Elements Have More than One Form • A few elements, including oxygen, phosphorus, sulfur, and carbon, are unusual because they exist as allotropes. • An allotropeis one of a number of different molecular forms of an element.

42. Compounds are Pure Substances • Pure substances that are not elements are compounds. • Compounds are composed of more than one kind of atom. • For example, carbon dioxide is a compound that is composed of molecules that consist of one atom of carbon and 2 atoms of oxygen. • Compounds are chemically bonded in specific ratios.

43. Compounds are Represented by Formulas • Because every molecule of a compound is made up of the same kinds of atoms arranged the same way, a compound has characteristic properties and composition. • For example, every molecule of hydrogen peroxide contains two atoms each of hydrogen and oxygen. This compound can be represented by an abbreviation or formula: H2O2 • Subscripts are placed to the lower right of the element’s symbol to show the number of atoms. • If there is just one atom, no subscript is used. • Molecular formulas give information only about what makes up a compound. • A structural formula shows how the atoms are connected, but the two-dimensional model does not show the molecule’s true shape.

44. Compounds Are Further Classified • By the type of bond that holds them together, and by whether they are made of certain elements.

45. Mixtures • A matter that contains two or more pure substances is a mixture. • The proportions of the ingredients can vary in a mixture.

46. Mixtures Can Vary in Composition and Properties • A mixture does not always have the same balance of ingredients. • Because of this the properties of mixtures may vary. • Alloy is a solid mixture.

47. Homogeneous Mixtures • The pure substances are distributed uniformly throughout the mixture. • Different components can not be seen, not even using a microscope. • Any 2 samples taken from the mixture will have the same proportions of ingredients.

48. Heterogeneous Mixtures • It contains substances that are not evenly mixed. • Different samples of the mixtures will have different properties.

49. Distinguishing mixtures from Compounds • A compound is composed of 2 or more elements chemically joined together. • Mixture: they are not chemically joined • Mixture: varying proportions.

50. Separating Mixtures • Chemist’s often handle the task of separating the components of a mixture based on one or more physical properties. • Techniques used by chemists include filtration, which relies on particle size and distillation and evaporation, which rely on differences in boiling point.