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Presentation on theme: "A Review of Chemistry Everything You Wish You’d Have Studied Harder Last Year! Or “Oh God, Not Again!” click on "Actions" button to download."— Presentation transcript:

1 A Review of Chemistry Everything You Wish You’d Have Studied Harder Last Year! Or “Oh God, Not Again!” click on "Actions" button to download

2 States of Matter Solids-the particles in a solid are packed tightly together so there is little movement. o definite volume-definite shape. Liquids-particles in a liquid are much farther apart and can move around more. definite volume-indefinite shape. o definite volume-indefinite shape. Gases-particles very far apart expand to fill container. o indefinite volume-indefinite shape.

3 Changes in State When a substance goes from one state to another. Solid  Liquid=melting (for H 2 O-32 o F or 0 o C) Liquid  gas=evaporating (for H 2 O-212 o F or 100 o C) Gas  liquid=condensing (for H 2 O-212 o F or 100 o C) Liquid  Solid=freezing (for H 2 O-32 o F or 0 o C)

4 Pure Substances and Mixtures Matter Pure SubstanceMixture ElementCompoundHomogeneousHeterogeneous A pure substances has a definite and constant composition. Pure substances can be elements like Gold or Potassium or compounds like salt or sugar. Elements are made of only 1 kind of atom. They cannot be taken apart. Compounds are made of 2 or more different kinds of atoms. They can be chemically taken apart. A mixture is a combination of pure substances. The composition may vary. Mixtures can be taken apart by filtration, distillation, evaporation etc. Homogeneous mixture is one that looks the same throughout like Koolaid. Heterogeneous mixture is one that is different throughout like a granola bar.

5 Measuring Matter The SI system o Decimal system with basic units for mass, length, and volume and prefixes that modify them. Units Prefixes  Mass=gram kilo (k) 1000  Length=meter centi (c).01  Volume=liter milli (m).001

6 Converting SI Units

7 Dimensional Analysis: SI/English Units Unit factors may be made from any two terms that describe the same or equivalent "amounts" of what we are interested in. For example, we know that 1 inch = 2.54 centimeters We can make two unit factors from this information:

8 Properties of Matter Chemical Properties -describe a substances ability to change into something else. o Flammability o Reactivity Physical Properties -describe a substances physical characteristics. o Mass o Color o Density

9 Physical Properties Extensive property- depends on how much matter is present. Mass and volume are extensive properties. Intensive properties-do not depend on amount of matter; a small chunk of gold is just as yellow as a large chunk of gold. Density d=m/v Density=mass */ * by volume Unit is g/cm 3 Measuring Density Irregular solid Water displacement Regular solid find mass, find volume divide

10 Temperature Scales

11 Converting Temperature from scale to scale K= o C + 273 o C=5/9 ( o F-32) o F=9/5 ( o C) + 32 What is normal body temperature 98.6 o F in o C? Answer 37 o C

12 Subatomic Particles Three major particles o Proton Eugen Goldstein cathode ray tube  +charge  1 amu (mass)  Located in nucleus o Neutron James Chadwick  no charge  1 amu (mass)  Located in nucleus o Electron J J Thomson cathode ray tube  -charge  No recognizable charge  Outside the nucleus

13 Dalton’s Atomic Theory 1. All elements are composed of tiny indivisible particles called atoms 2. Atoms of the same element are identical. The atoms of any one element are different from those of any other element 3. Atoms of different elements can physically mix together or can chemically combine in simple whole- number ratios to form compounds 4. Chemical reactions occur when atoms are separated, joined or rearranged. Atoms of one element are never changed into atoms of another element as the result of a chemical change

14 Two Changes to Dalton’s Theory 1. Atoms are not indivisible Rutherford’s Nuclear Model o They have “stuff” inside  Nucleus  Protons  Neutrons  Quarks  Neutrinos 2. Atoms of one element can be changed to different element during a nuclear reaction.

15 The Nucleus 1911 Ernest Rutherford discovered the nucleus with the gold foil experiment. Very small!!! If Superdome in New Orleans represented an atom, the nucleus would be about the size of a pea on the 50 yard line. Holds protons and neutrons, thus most all of the mass of an atom.

16 Protons + Neutrons Sum of protons + neutrons = mass number. Number of protons =atomic number

17 Calculating Protons, Neutrons and Electrons Uranium’s atomic number is 92 (# or protons) Uranium’s mass number is 238 (protons plus neutrons) To find # of neutrons in U, subtract the atomic # from the mass #. # neutrons =146 # of electrons = mass # (In a neutral atom)

18 Electrons Electrons can be considered to orbit the nucleus in certain energy levels. The outermost electrons from the nucleus are called the valence electrons. Valence electrons are involved in bonding. Number of valence electrons also determine atoms reactivity.

19 Isotopes An isotope is an atom that has the same number of protons (or the same atomic number) but different number of neutrons (or different atomic mass).

20 Ions In neutral atoms, the number of electrons = the number of protons. Ions have a different number of electrons. Some atoms gain electrons and some lose electrons. Atoms that gain e- become negatively charged ions or anions. Atoms that lose e- become positively charged ions or cations.

21 The Periodic Table Mid-1800’s Dmitri Mendeleev noticed a repeating pattern of chemical properties in the elements that were known at the time. Mendeleev first arranged the elements in order of increasing atomic mass. Very similar to todays PT. Note exceptions---Co-Ni, Te-O

22 Arrangement of the PT Elements arranged in increasing atomic number. Horizontal rows 1-7 are periods. o All elements in a period have electrons in same energy level. Vertical columns 1-18 are groups or families. o Members of families have same number of valence electrons thus, similar properties Metals left of staircase, nonmetals right of staircase, metalloids touching the staircase.

23 Families or Groups Have similar properties IA-Alkali Metals o Li, Na, K, Rb, Cs, Fr o Lose one electron in reactions o React violently in water o Commonly found in salts o Important in bodily chemistry

24 Families or Groups cont. IIA-Alkaline earth metals o Be, Mg, Ca, Sr, Ba, Ra o Lose two electrons in reactions VIIA-Halogens o F, Cl, Br, I o Gain one electron VIIIA-Noble Gases o He, Ne, Ar, Kr, Xe, Rn o Very unreactive

25 A Groups Number in front of A Group Columns (1A, 2A, 3A, 4A, 5A, 6A, 7A, 8A) corresponds to the number of valence electrons. o 1A-all have one valence e- o 2A-all have two valence e- o 7A-all have seven valence e- o 8A-have eight valence e--the magic number! The octet makes the atom stable!

26 Periodic Trends

27 Metals Solid (except for Hg, a liquid) Shiny Good conductors of electricity and heat Ductile (can be drawn into thin wire) Malleable (easily hammered into thin sheets) Tend to lose electrons to become cations Majority of elements are classified as metals

28 Nonmetals Brittle Not malleable Not ductile Poor conductors of heat and electricity Gain electrons to become anions Some nonmetals are liquid

29 Metalloids: semimetals Partially conduct electricity o Semiconductors o Computer chips

30 Bonding: Ionic Bonds Ionic bonds are formed between a metal and a nonmetal. The metal wants to lose electrons and the nonmetal wants to gain electrons to achieve a stable octet of valence electrons. The transfer of electrons creates ions with opposite charges. Opposite charges attract each other.

31 The Octet Rule Atoms will gain, lose, or share electrons until there is a total of 8 valence electrons in their outer most energy level. 1A-have one too many, easier to lose 1 electron than gain 7. 7A-need only 1 more to have 8, easier to gain 1 than to lose 7.

32 Ionic charge You can often determine the ions charge by it’s position on the periodic table. All group 1A metals have a +1 charge. Group 2A metals have a +2 charge Group 3A metals have a +3 charge. Group 7A nonmetals have a -1 charge Group 6A nonmetals have a -2charge Group 5A nonmetals have a -3 charge

33 Transition metals Group/Family B- cannot determine charge by group #. Must memorize charges. o Fe +2, Fe +3 o Cu +1, Cu +2 o Mn +2, Mn +3 o Co +2, Co +3 o Sn +2, Sn +3 o Pb +2, Pb +4

34 Polyatomic Ions Ions aren’t always composed of only one kind of atom; sometimes they are polyatomic, composed of more than one kind of atom. Sulfate SO 4 -2 1 sulfur and 4 oxygen Nitrate NO 3 -1 1 nitrogen and 3 oxygen Chlorate ClO 3 -1 1 chlorine and 3 oxygen Phosphate PO 4 -3 1 phosphate and 4 oxygen Carbonate CO 3 -2 1 carbon and 3 oxygen Hydroxide OH -1 1 oxygen and 1 hydrogen Ammonium NH 4 +1 1 nitrogen and 4 hydrogne

35 Ionic Compounds When ionic compound is formed the cation and the anion attract each other-compound must be neutral! Writing ionic compound formulas Put metal and nonmetal side by side. Metal first!  Mg +2 Br -1  If charges add up to zero then just write the formula.  If not, crisscross the charges to balance the formula.  Charge on opposite ion become the subscript on other ion.  MgBr 2

36 Bonding: Covalent Bonds A covalent bond is a bond between two or more nonmetals. In a covalent bond, electrons are shared, not transferred. No ions! Same goal- octet of valence electrons. Covalent compounds may share more than one bond; two or three! (double or triple bond) more than one pair of electrons are shared

37 Covalent Bonding Diatomic molecules 7 naturally occuring diatomic molecules H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2 Lewis Dot Structures o Valence electrons are represented as dots surrounding the atomic symbol.

38 Exceptions to the Octet Rule NO 2, PCl 5, SF 6, BF 3 The octet rule cannot be satisfied in molecules whose total number of valence electrons is an odd number.

39 VSEPR Theory Atoms in a compound will repel each other Because of this, atoms within a compound will situate themselves so that they are as far apart as possible.

40 Naming Covalent Bonds Add prefix Mono =1 Di=2 Tri=3 Tetra=4 Penta=5 Hexa=6 Hepta=7 Octa=8 Nona=9 Deca=10

41 Chemical Reactions Exothermic o Gives off heat Endothermic o Takes heat in

42 Types of Reactions Combination/Synthesis Decomposition Single Displacement Double Displacement Combustion

43 Balancing Reactions The Law of Conservation of Mass states that matter is neither created or destroyed in a chemical reaction. Thus you must have the same amout of stuff in your product as in the reactant.

44 Mole Road Map The road map can be used a guideline to tell you how to use the mole in converting information. **Just decide what you are starting with and what you want to figure out. Then find those two points on the map above and follow the roads. **The roads represent the steps you will have to do and the conversion factors you will need to know in order to accomplish what you want.

45

46 Molar Mass Add the atomic masses of all atoms in a compound NH 3 H 2 SO 4 CaCl 2

47 %Composition Part/Whole x 100 KCl NaF

48 Molarity What is the concentration of the solution? M=m/L

49 Dilutions M 1 V 1 =M 2 V 2

50 Acid/Base Acid Ph of less than 7 Have H+in the front of the formula Strong acids decompose completely in water Bases Ph of greater than 7 Usually contain OH- Strong bases decompose completely in water

51 Gas Laws 4 Important physical properties of gases o Volume o Pressure o Temperature o Amount-moles

52 Boyles Law Pressure-Volume relationship if temp and amount stays the same P 1 V 1 =P 2 V 2

53 Charles Law Temp-volume relationship if amount and pressure are kept the same V 1 /T 1 =V 2 /T 2

54 Gay-Lussac’s Law Pressure-temp relationship when amount and volume are kept same P 1 /T 1 =P 2 /T 2

55 Combined Gas Law When the properties cannot be held constant. P 1 V 1 /T 1 =P 2 V 2 /T 2

56 Stoichiometry: How much needed? How much made? Mole Ratio


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