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Date Created: 06/03/16
College Chemistry 1 – Mona Chaudhary 5/23/2016 Important Words or Important People Important Concepts Definitions Lesson 2 Measurements and Units English & Metric oz, lb, ton Gram, Kg, mg in, ft, yd, mile cm, meter, Km pint, quart, gallon Liter, mL Significant Figures o Non zero digits 38.56 (4) 288 (3) o Zeros at the beginning of a number is not significant and are considered place holders 0.052 (2) 0.00321 (3) o Zeros between non zeros are significant 2007 (4) 302 (3) o Zero at the end of a decimal number are significant 38.0 (3) o Zero at the end of a non-decimal number are ambiguous 24,300 (3, because there are no units) 24,300 Km (5, because there is a unit) o A decimal at the end of the number makes zeros significant 280. (3) 2. Ft (1) Rounding o < 5, leave it 6.474 = 6.47 o > 5, round up 6.5483 = 6.55 o = 5, round up 6.575 = 6.58 Dimensional analysis o Converting from one unit to another using factors o How many inches are in 3.2 miles? (5280 ft = 1 mile) Significant figures are based on the least amount of sig. figs. in the question. In this case, there are 3 3.2 Miles 5280 ft 12 inches 20,2752 inches 5 X 1 Mile X 1 ft = 2.03 x 10 inches X Density o D = M/V o V = M/D o M = VD Temperature o Fahrenheit - °F (1.8C)+32 o Celsius - °C (F-32)/1.8 o Kelvin - K C+273 College Chemistry 1 – Mona Chaudhary 5/25/2016 Important Words or Important People Important Concepts Definitions Lesson 3 Matter o Has weight, shape, and mass Desk, classroom Composed of atoms Atom o Smallest unit that can be identified as a particular element o Fundamental Particles (building blocks) Proton P or p+ Mass ≈ 1 amu Charge = +1 Location = in the nucleus Neutron n or n0 Mass ≈ 1 amu Charge = +1 Location = in the nucleus Electron e or e Mass ≈ 0 Charge = -1 Location = outside the nucleus o Concentric circles Atomic Number (Z) o Always a whole number o The number of protons in the nucleus + o Z = # p o Identifies element in the periodic table Hydrogen Z = 1 Helium Z = 2 Z = 17? Chlorine (Cl) Z = 76? Osmium (Os) o Neutral atoms The number of protons = the number of electrons + - H is different than H or H Cl Cl- H H+ Li2 + + + + + p = 17 p = 17 p = 1 p = 1 p = 3 e = 17 e = 18 e = 1 e = 10 e = 1 ANION CATION CATION Mass Number o Always a whole number + 0 o A = p + 0 o A-Z = n Isotopes: Isotopes o Atoms of the same element with different masses p = 29 p = 29 o All isotopes have the same number of protons (Z)63 65 29 Cu e = 29 29 Cu e = 29 o May have a different number of neutrons n = 34 n = 36 Atomic Weight / Mass o ALWAYS a decimal number o Weighted average of all the isotopes of a particular element o Round to two places after the decimal Ions o Charges particles formed from a gain or loss of electrons The number of protons stays the same o Cation – positive charge – less electrons than protons o Anion – negative charge – more electrons than protons College Chemistry 1 – Mona Chaudhary 5/25/2016 Important Words or Important People Important Concepts Definitions Lesson 4 Electron Configuration Types o Bohr – electrons are fixed in concentric, circular orbits at a specified distance from the nucleus o Quantum Mechanical – electrons are defined to regions with in energy levels outside the nucleus Electron Configuration o Describes the arrangement of electrons o Methods: Shorthand Notation Orbital Energy Diagram They both give the same information o Components: Principal energy level (1-7) S = electron Orbital holding e (s,p,f,d) Subscript on orbital letter = #e - P present o Orbital Notes S can have 1-2 electrons D P can have 1-6 D can have 1-10 F F can have 1-14 Shorthand Notation o Level, orbital, e- o The periodic table gives all 3 components Level = period 1-7 Orbital = Groups I & II = S III – VIII = P Transition Metals = D # e = number of electrons in outer shell (+/- for ions) o EXAMPLE H = 1s 1 He = 1s 2 Orbital Energy Diagrams o Graphic representation of shorthand notation Written from bottom to top Orbitals are represented by dashed line(s) s orbital = 1 dash p orbital = 3 dashes d orbital = 5 dashes f orbital = 7 dashes (we do not show) Aufbau Principle – fill the lowest energy level first Hund’s Rule – Fill in orbitals unpaired first Give each orbital an electron first, then go back and double up Draw arrows in the opposite direction because electrons are magnetic and opposite ends attract Octet configurations are extremely stable because the valence (outermost) orbital is filled with electrons. o Duet configurations are also extremely stable Examples: 1 1H= 1s 1s 32Ge = 1s , 2s , 2p , 3s , 3p , 4s , 3d , 4p 2 4p 3d 4s 3p 3s 2p 2s 1s College Chemistry 1 – Mona Chaudhary 6/1/2016 Important Words or Definitions Important People Important Concepts Lesson 5 Periodic Table – an organized listing of all known types of atoms with some important information for each atom or element Atomic Number 1 1.009 Atomic Mass (weight) Chemical Symbol H Name of Hydrogen element o Atomic Number – The number of protons found in the nucleus o Never Changes o Atomic Mass – the number of protons plus the number A of neutrons found in the nucleus. X Z o This is the average number, because the number of X = element neutrons can change. o Isotopes – atoms of the same element with the same A = p + n 0 number of protons, but different numbers of neutrons + (therefor having a Z = p different atomic 12 13 14 weight) 6C 6C 6C Calculations: o The atomic mass – the atomic number equals the number of neutrons o Atoms usually have a zero net charge o Because the protons and electrons are the same amount (balanced) Na Groups ↕ Atomic Number = 11 o Main groups Mass = 23 Main group metals Protons = 11 o Alkali Metals Neutrons = 12 Electrons = 11 1A EXCEPT Hydrogen o Alkaline Earth Metals 2A o Halogens 7A o Noble Gases (Inert Gases) 8A o Transition Metals / “d” block metals 3B-2B o Metals, Metalloids, Non-metals Characteristics Metals – conductive, shiny, has a shape Non-Metals – dull, not conductive Metalloids – has characteristics of both o Inner Transition Elements Lanthanide series Actinide series Periods ↔ o 7 o =Energy Levels Trends o Atomic Radius ← ↓ o Electronegativity ↑ → o Ionization Energy ↑ → o Reactivity Metals – Large atoms are more reactive ← ↓ Non-metals – smaller atoms are more reactive ↑ → College Chemistry 1 – Mona Chaudhary 6/1/2016 Important Words or Definitions Important People Important Concepts Lesson 6 Bonds o Ionic – an attraction between negatively and e positively charged ions - ionic bonds form between: P P Metals & Non-Metals e strong electron donors - o Metals Hydrogen Gas o 1 or 2 electrons in valence shell Strong electron acceptor o Non-Metals o 6 or 7 electrons in valence shell o Covalent – when two orbitals overlap and electrons are shared “Share a pair” Non-Metals & Non-Metals A single line drawn between two atoms represents the sharing of 2 electrons H – H o 2 electrons / 1 pair O=O o Double covalent bond o 4 electrons / 2 pair Nonpolar covalent bond – when the sharing of electrons is fair and equal Methane, hydrogen gas, oxygen gas Electronegativity – the attraction of an atom for electrons Polar covalent bond – unequal sharing of electrons small ionization potential o willingness to lose or give up electrons Positive end Large value for the electron affinity o More likely to grab an electron Negative end Linus Pauling – each atom has an ability to attract electrons – this ability is called electronegativity o Electronegativity change is greater than 0, less than 1.9 o The element with the higher electronegativity value will win the electron tug-of-war o Bottom left to upper right increases electronegativity EXCLUDING noble gases (they are already happy with the amount of electrons they have) Metals – low electronegativity Non-metals – high electronegativity Water O ẟ- o Polar molecule o Oxygen is more electronegative than the ẟ+ ẟ+ H H hydrogens Therefor it has a stronger pull on the electrons Lewis Dot Structures o Shows valence electrons Group number = #e - Group V = 5 valence electrons Drawing Lewis Dot Structures o Great for showing bonding Which atoms are bonded to which Shows the number and types of bonds o Not good for showing shape o Octet Rule – each atom wants 8 electrons Hydrogen is the only atom that only wants 2 (Duet rule) o NASB Method N = Needed Needed Available o Add: o 8 for every atom Shared o Only 2 for hydrogen Bonds o Always an even # A = Available - single = double o Add the number of valence electrons = triple S = Shared o N-A o Central atom is one of the following: B = Bonds The least electronegative Carbon o S/2 The only one of its kind o Not always correct Never Hydrogen o If drawing an ion, put the structure in square brackets [ ] NASB Practice: 1. PBr3 2. N 2 2 3. CH OH 3 4. C 2 4 College Chemistry 1 – Mona Chaudhary 6/6/2016 Important Words or Definitions Important People Important Concepts Lesson 7 Bonds o Ionic bonding A transfer of electrons Electronegativity difference greater than 1.9 Between metals and non-metals The farther apart they are, the more likely they are to bond o Covalent Sharing of electrons “Share a pair” Electronegativity difference between 0 and 1.9 Between nonmetals and non-metals Non-Polar / Pure When the difference in electronegativity is 0 Equal sharing of electrons Diatomic molecules o H O2N 2l 2 2 o Polar When the difference in electronegativity is greater than 0 but less than 1.9 Unequal sharing of electrons Partial charges represented by delta (ẟ+ or ẟ-) Dipole NO, CO, H O 2 Intermolecular forces of attraction o Attraction between molecules o Forces hold molecules close to each other London/Dispersion force Weakest Non-polar molecules Dipole-Dipole Intermediate force Polar molecules Hydrogen Bonding Strongest Only when hydrogen bonds with N, O, and F o Hydrogen bonding is “FON” College Chemistry 1 – Mona Chaudhary 6/6/2016 Important Words or Definitions Important People Important Concepts Lesson 8 Naming Compounds Greek Prefix: What is the first element? Non-Metal 1 = mono Non-Metal + anion 2 = di Use the Greek prefix before each name to show how many there are 3 = tri o CO /2Carbon dioxide 4 = tetra Group I, II, or III metal 5 = penta Name the metal and anion o KO / Potassium Oxide 6 = hexa Transition metal (or Sb, Sn, Bi, or Pb) 7 = hepta Determine the charge of the transition metal o You can usually tell the charge by the 8 = octa subscript on the anion 9 = nona o If not, you have to calculate the charge 10 = deca of the anion, and divide by the number of transition metal atoms (the subscript) Metal (charge in roman numerals) + anion Writing Formulas o Write down the first element (or Poly Atomic Ion) o If it’s a metal: Roman Numerals: Group I, II, or III Write the atom and charge (group number) 1 = I Write the anion and charge 2 = II Determine the number of each atom needed 3 = III to make the total charge neutral (show as subscript) 4 = IV Transition metal (or Sb, Sn, Bi, or Pb) 5 = V Write the atom and charge (whatever is in Roman Numerals) 6 = VI Write the anion and charge 7 = VII Determine the number of each atom needed 8 = VIII to make the total charge neutral (show as subscript) College Chemistry 1 – Mona Chaudhary Unit 1 Study Guide Concepts to Review Measurements and Units Define Cont: English & Metric o Neutron o Electron oz, lb, ton Gram, Kg, mg Atomic Number in, ft, yd, mile cm, meter, Km Mass Number Isotopes pint, quart, gallon Liter, mL Atomic Weight/Mass Significant Figures Ion Non zero digits o Cation Zeros at the beginning of a number o Anion Zeros between non zeros Concepts: Zero at the end of a decimal number Shorthand vs longhand electron Zero at the end of a non-decimal number configuration notation A decimal at the end of the number Bohr and Quantum Mechanical electron configurations Rounding Orbital Energy Diagrams Periodic Table < 5, leave it > 5, round up o Groups = 5, round up Main groups Alkali Metals Dimensional analysis Alkaline Earth Metals Know how to use conversion units Halogens Density Noble Gases Transition Metals D = M/V Inner Transition Elements Lanthanide Series Temperature Actinide Series Fahrenheit - °F o Periods o (1.8C)+32 o Trends Celsius - °C Atomic Radius o (F-32)/1.8 Electronegativity Kelvin - K Ionization o C+273 Reactivity People to Know: Metals and Non- Metals Bohr Lewis Dot Structures Aufbau o NASB Method Hund Intermolecular Forces of Attraction Define: o London / Dispersion o Dipole – Dipole Matter Atom o Hydrogen Bonding Naming Compounds given formulas o Proton Writing Formulas given compounds Sample Questions: Write the name and symbol for the neutral element bellow: 1. Atomic mass = 196.967 amu 2. Electronic Configuration is: 1s , 2s , 2p , 3s , 3p , 4s , 3d , 4p 2 6 3. Electronic Configuration is: [Kr] 5s , 4d 4. Has 37 protons: 5. Neutral atom has 17 electrons: 6. Mass number is 207 and there are 125 neutrons 7. Forms an anion that has a 2- charge and 50 electrons: 8. Complete the table: Atom/Ion Atomic Mass Number Protons Neutrons Electrons Number Hg 30 31 56 9. Write the longhand electronic configuration for P 3- 10. Write the longhand electronic configuration for Br 11. Draw, label, and fill the orbital energy diagram for Ga: 12. Arrange the elements from smallest to largest according to atomic radius (Ni, Ar, La, Li) 13. Arrange the elements from largest to smallest according to the amount of electronegativity (Ge, Sb, He, Mo) 14. Draw a Lewis Dot Structure and show the NASB method (show your work) CH OH 3 15. Round the number 3785.642 to 3 significant numbers 16. Put 2319 into scientific notation 17. Group _____ contains Alkaline Earth Metals a. 1A b. 2A c. 6A d. 7A 18. He, Ne, Ar, Kr, Xe, and Rn are the _____________ __________. a. Alkali Metals b. Alkaline Earth Metals c. Halogens d. Noble Gases 19. Chlorine has _____ valence electrons. a. 3 b. 7 c. 8 d. 5 20. There are _____ significant figures in : 308.640 a. 5 b. 4 c. 6 d. 7 21. Arrange the following elements from lowest to highest in ionization energy: S, Ta, Sc, Fr 22. Calculate the number of inches in 3.7 Kilometers. (39.37 inches / meter) Show work 23. Write the names for the following: a. P4 7 b. CaBr 2 c. Ba(NO2) 2 d. Fe2O 3 e. Sn3(PO 4 4 24. Write the formula for the following: a. dicarbon hexachloride b. bismuth (V) phosphate c. aluminum chlorite d. lithium fluoride e. iron (II) perchlorate
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