CHEMISTRY, Julia Burge 2 nd edition. Copyright McGraw-Hill

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1 Chapter 1 Chemistry: the Central Science CHEMISTRY, Julia Burge 2 nd edition Copyright McGraw-Hill

2 11The 1.1 Study of Chemistry Chemistry the study of matter and the changes that matter undergoes Matter anything that has mass and occupies space

3 13Classification 1.3 of Matter Matter is either classified as a substance or a mixture of substances Substance Can be either an element or a compound Has a definite (constant) composition and distinct properties Examples: sodium chloride, water, oxygen

4 States of Matter Solid particles close together in orderly fashion little freedom of motion a solid sample does not conform to the shape of its container Liquid id particles close together but not held rigidly in position particles are free to move past one another a liquid sample conforms to the shape of the part of the container it fills

5 Gas particles randomly spread apart particles have complete freedom of movement a gas sample assumes both shape and volume of container. States of matter can be inter converted without changing chemical composition solid liquid gas (add heat) gas liquid solid (remove heat)

6 States of Matter

7 Substances Element: cannot be separated into simpler substances by chemical means. Examples: iron, mercury, oxygen, and hydrogen Compounds: two or more elements chemically combined in definite i ratios Cannot be separated by physical means Examples: salt, water and carbon dioxide

8 Mixtures Mixture: physical combination of two or more substances Substances retain distinct identities No universal constant composition Can be separated by physical means Examples: sugar/iron; sugar/water

9 Molecular Comparison of Substances and Mixtures Atoms of an element Molecules l of an element Molecules of a compound Mixture of two elements and a compound

10 Types of Mixtures Homogeneous: composition of the mixture is uniform throughout Example: sugar dissolved in water Heterogeneous: composition is not uniform throughout Example: sugar mixed with iron filings

11 Classification of Matter

12 Classify the following Aluminum foil: Baking soda: Milk: Air: Copper wire: substance, element substance, compound mixture, homogeneous mixture, homogeneous substance, element

13 1.3 Scientific Measurement Used to measure quantitative properties of matter SI (System International) base units

14 Why are units important? Mars Climate Orbiter ($125 million) destroyed by wrong conversion units

15 SI Prefixes

16 Mass: measure of the amount of matter (weighti h refers to gravitational i pull) Temperature: Celsius Represented by C Based on freezing gpoint of water as 0 C and boiling point of water as 100 C Kelvin Represented dby K( (no degree sign) The absolute scale Units of Celsius and Kelvin are equal in magnitude Fahrenheit(the English system) ( F)

17 Equations for Temperature Conversions o C = ( o F 32) 5 9 K= o C o 9 F = o C

18 Temperature Conversions A clock on a local bank reported a temperature t reading of 28 o CWhti C. What is this temperature on the Kelvin scale? K = o C K = 28 o C = 301K

19 Practice Convert the temperature reading on the local bank (28 C) into the corresponding Fahrenheit temperature. o F 9 = 5 o C + 32 o F = o C + 32 = 82 o F

20 Volume: meter cubed (m 3 ) Derived unit The unit liter (L) is more commonly used in the laboratory setting. It is equal to a decimeter cubed (dm 3 ).

21 Density: Ratio of mass to volume Formula: d = m V d = density (g/ml) m = mass (g) V = volume (ml or cm 3 ) (*gas (gas densities are usually expressed in g/l)

22 Practice The density of a piece of copper wire is 8.96 g/cm 3. Calculate the volume in cm 3 of a piece of copper with a mass of 4.28 g. m d = V V m = d = 4.28 g = g 8.96 cm cm 3

23 14Properties 1.4 of Matter Quantitative: expressed using numbers Qualitative: expressed using properties Physical properties: can be observed and measured without changing the substance Examples: color, melting point, states of matter Physical changes: the identity of the substance stays the same Examples: changes of state (melting, freezing)

24 Chemical properties: must be determined by the chemical changes that are observed Examples: flammability, acidity, corrosiveness, reactivity Chemical changes: after a chemical change, the original substance no longer exists Examples: combustion, digestion

25 15Uncertainty 1.5 in Measurement Exact: numbers with defined values Examples: counting numbers, conversion factors based on definitions Inexact: numbers obtained by any method other than counting Examples: measured values in the laboratory

26 Practice L L = L Calculator l answer: L Round to: L Answer to the tenth position cm x cm x cm = cm 3 Calculator answer: cm 3 Round to: cm 3 round to the smallest number of significant figures

27 Accuracy and Precision Accuracy and precision Two ways to gauge the quality of a set of measured numbers Accuracy: how close a measurement is to the true or accepted value Precision: how closely measurements of the same thing are to one another

28 both accurate and precise not accurate but precise neither accurate nor precise

29 Describe accuracy and precision for each set Student A Student B Student C g g g g g g g g g Average: g g g True mass is grams

30 Student A s results are precise but not accurate. Student B s results are neither precise nor accurate. Student C s Cs results areboth precise and accurate.

31 1.6 Using Units and Solving Problems Conversion factor: a fraction in which the same quantity is expressed one way in the numerator and another way in the denominator Example: by definition, 1 inch = 2.54 cm 1in 2.54 cm 2.54 cm 1in

32 Dimensional analysis: a problem solving method employing conversion factors to change one measure e to another often called the factor label method Example: Convert inches to meters Conversion factors needed: 2.54 cm = 1 in and 100 cm = 1 meter cm 1m in = 1in 100 cm m

33 Notes on Problem Solving Read carefully; find information given and what is asked for Find appropriate equations, constants, conversion factors Check for sign, units and significant figures Check for reasonable answer

34 Practice The Food and Drug Administration (FDA) recommends that t dietary sodium intake be no more than 2400 mg per day. What is this mass in pounds (lb), if 1 lb = g? 2400 mg 1g 1000 mg 1lb g = lb

35 Scientific method Key Points Classifying matter Density Temperature conversions Physical vs chemical properties and changes Precision vs accuracy Dimensional analysis

36 QUESTIONS? 36

37 Chapter 2 Atoms, Molecules, and Ions 37

38 2.3 Atomic Number, Mass Number and Isotopes The chemical identity of an atom can be determined solely from its atomic number Atomic number (Z) number of protons in the nucleus of each atom of an element Also indicates number of electrons in the atom since atoms are neutral 38

39 Mass number (A) total number of neutrons and protons present in the nucleus Standard notation: 39

40 Isotopes All atoms are not identical (as had been proposed by Dalton) Same atomic number (Z) but different mass numbers (A) Isotopes of Hydrogen Hydrogen (protium) Deuterium Tritium 40

41 24The 2.4 Periodic Table Periods horizontal rows Families (Groups) vertical columns Elements in the same family have similar chemical and physical properties Arranged in order of increasing atomic number 41

42 The Modern Periodic Table Ti 42

43 Metals good conductors of heat and electricity (majority of elements on the table, located to the left of the stair step) Nonmetals nonconductors (located in upper right hand corner) Metalloids in between metals and nonmetals (those that lie along the separation line) 43

44 Groups (Families) on the Periodic Table 44

45 2.5 The Atomic Mass Scale and Average atomic Mass Atomic mass is the mass of the atom in atomic mass units (amu) Atomic mass unit is defined as a mass exactly equal to one twelfth the mass of one carbon 12 atom Carbon 12 (12 amu) provides the standard for measuring the atomic mass of the other elements 45

46 2.6 Molecules and Molecular Compounds Molecule combination of at least two atoms in a specific arrangement held together by chemical bonds May be an element or a compound H 2, hydrogen gas, is an element H 2 O, water, is a compound 46

47 Diatomic molecules: Homonuclear (2 of the same atoms) Examples: H 2 2, N 2 2, O 2 2, F 2 2, Cl 2 2, Br 2 2, and I 2 Heteronuclear (2 different atoms) Examples: CO and HCl 47

48 Polyatomic molecules: Contain more than 2 atoms Most molecules May contain more than one element Examples: ozone, O 3 ; white phosphorus, P 4 ; water, H 2 O, and methane (CH 4 ) 48

49 Molecular formula shows exact number of atoms of each element in a molecule Subscripts indicate number of atoms of each element present in the formula. Example: C 12 H 22 O 11 Structural formula shows the general arrangement of atoms within the molecule. 49

50 Naming molecular compounds Binary Molecular compounds Composed of two nonmetals Name the first element Name the second element changing ending to ide Use prefixes to indicate number of atoms of each element 50

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53 Name the following: NO 2 nitrogen dioxide N 2 O 4 dinitrogen tetraoxide Write formulas for the following: Diphosphorus pentoxide P 2 O 5 Sulfur hexafluoride SF 6 53

54 Common Names B 2 H 6 diborane SiH 4 silane NH 3 ammonia PH 3 phosphine hi H 2 O water H 2 S hydrogen sulfide 54

55 Acid a substance that produces hydrogen ions (H + ) when dissolved in water Binary acids: Many have 2 names Pure substance Aqueous solution Example: HCl, hydrogen chloride, when dissolved in water it is called hydrochloric acid 55

56 Naming binary acids Remove the gen ending from hydrogen (leaving hydro ) Change the ide ending on the second element to ic Combine the two words and add the word acid. 56

57 Name the following: HBr hydrogen bromide hydrobromic acid H 2 S hydrogen sulfide hydrosulfuric acid Write formulas for the following: Hydrochloric acid HCl(aq) Hydrofluoric acid HF(aq) 57

58 Organic compounds contain carbon and hydrogen (sometimes with oxygen, nitrogen, sulfur and the halogens.) Hydrocarbons contain only carbon and hydrogen Alkanes simplest examples of hydrocarbons Many derivatives of alkanes are derived by replacing a hydrogen with one of the functional groups. Functional group determines chemical properties 58

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61 Empirical formulas reveal the elements present and in what whole number ratio they are combined. Molecular(explicit) Empirical(simplest) H 2 O 2 HO N 2 H 4 NH 2 H 2 O H 2 O 61

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63 2.7 Ions and Ionic Compounds Ion an atom or group of atoms that has a net positive or negative charge Monatomic ion one atom with a positive or negative charge Cation ion with a net positive charge due to the loss of one or more electrons Anion ion with a net negative charge due the gain of one or more electrons 63

64 Common Monatomic Ions 64

65 Naming ions Cations from A group metals Name the element and add the word ion Example: Na +, sodium ion Cations from transition metals with some exceptions Name element Indicate charge of metal with Roman numeral Add word ion Example: Cu 2+,copper(II) ion 65

66 Anions Name the element and modify the ending to ide Example: Cl, chloride Polyatomic ions ions that are a combination of two or more atoms Notice similarities number of oxygen atoms and endings for oxoanions Nitrate, NO 3 and nitrite, NO 2 66

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68 Ionic compounds represented by empirical formulas Compound formed is electrically neutral Sum of the charges on the cation(s) and anion(s) in each formula unit must be zero Examples: Al 3+ and O 2 Al 2 O 3 Ca 2+ and PO 4 3 Ca 3 (PO 4 ) 2 68

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70 Formation of an Ionic Compound 70

71 Write empirical formulas for Aluminum and bromide AlBr 3 barium and phosphate Ba 3 (PO 4 ) 2 Magnesium and nitrate Mg(NO 3 ) 2 Ammonium and sulfate (NH 4 ) 2 SO 4 71

72 Naming ionic compounds Name the cation Name the anion Check the name of cation If it is a A group metal you are finished If it is a transition metal, with some exceptions, add the appropriate Roman numeral to indicate the positive ionic charge 72

73 Write names for the following: KMnO 4 potassium permanganate Sr 3 (PO 4 ) 2 strontium phosphate Co(NO 3 ) 2 cobalt(ii) nitrate FeSO 4 iron(ii) sulfate 73

74 Hydrates compounds that have a specific number of water molecules within their solid structure Hydrated compounds may be heated to remove the water forming an anhydrous compound Name the compound and add the word hydrate. Indicate the number of water molecules with a prefix on hydrate. Example: CuSO 4 5H 2 O Copper (II) sulfate pentahydrate See picture 74

75 CuSO 4 CuSO 4.5H 2 O 75

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78 Key Points Isotopes Periodic table; families and periods; metals, nonmetals and metalloids Average atomic mass Naming and writing formulas for Binary molecular compounds Binary acids Ionic compounds Hydrates 78

79 QUESTIONS? 79