Cheat Sheet for Chemistry PDF

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Crib sheet for Final Exam...

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Chemistry Midterm The study of the composition, properties, and transformation of matter

Mole Calculations: Multiplication or division; fewest SF as the smallest measurement. Addition or subtraction: fewest decimal places as the smaller measurement

Sig-Figs: Multiplication or division; fewest SF as the smallest measurement. Addition or subtraction: fewest decimal places as the smaller measurement

1.1 Operate units of measurement. - Percentage concentration o By mass (equation) o By volume (equation) - Molarity (mol/L): number of moles of a substance dissolved in each liter of solution - Solutions: a solute uniformly dispersed in a solvent o Precipitation reaction: forming an insoluble solid called a precipitate o Acid-base neutralization reaction: forming water o Oxidation-reduction reaction: electron transfer occurring 1.3 Recognize experimentation and measurement. 2.0 Nuclear Structure of Atoms 2.1 Describe the atomic theory of matter. - John Dalton atomic theory 1. All matter is composed of atoms 2. Atoms of one element differ from those of other elements 3. Atoms combine in simple, fixed, whole-number ratios 4. Atoms are not created or destroyed in chemical reactions 2.2 Describe the properties of subatomic particles. 2.3 Explain isotopes of an element and how its atomic mass is calculated. (pictures 1 & 2) - Isotopes: atoms with identical atomic numbers but different mass number (same # of protons, different # of neutrons) - Mass number (A) = number of protons (Z) + number of neutrons (N) - Atomic number (Z) = number of protons - Atomic mass (amu): 1 amu = 1.6606x10-24 g 3.0 Electronic Structure of Atoms and Periodic Table - Structure of Atoms o Atom diameter: ~10-10 m o Subatomic particles  Protons (p or Z): +1 charge (9.109382 x10-28 g)  Neutrons (N): 0 charge (1.672622 x10-24 g)  Electrons (e-): -1 charge (1.674927 x10-24 g) 

Neutrons are almost identical in mass to protons which is similar to a hydrogen atom  The mass is primarily in the nucleus Forces: protons attract electrons; protons repel protons; electrons repel electrons  Protons stay together because of nuclear force  4 forces of nature: electromagnetic, weak nuclear, gravitational and nuclear forces 

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Structure of the Periodic Table (include small table outline if there is room) o 7 periods and 18 groups

3.1 Describe the arrangement of elements in the periodic table. - Mendeleev arranged the elements according to atomic weight and chemical reactivity - Arranged by increasing atomic weight - Atomic radius o Increases down a group; decreases left to right - Ionization energy o Energy required to remove an electron from an atom o Decreases down a group, increase left to right - Electronegativity o Ability of atom in a molecule to attract the shared electrons in a covalent bond o Decreases down a group; increases left to right 3.2 Describe the Bohr Model, concept of shells and sub-shells, and atomic line spectrum. - Bohr model picture - Describes the atom as a central nucleus containing protons and neutrons being orbited by electrons in shells - Balmer-Rydberg: equation o m: shell the transition is to (inner-shell) o n: shell the transition is from (outer-shell) - Quantum mechanics o Diagram o Electron spin: o Pauli exclusion: no two electrons in an atom can have the same four quantum numbers - Orbital-filling diagrams o Lower-energy are filled before higherenergy orbitals o One electron goes into each until all are half full

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Electron (↑) and electron pairs (↑↓)

3.3 Determine the number of valence electrons of atoms. - Valence shell: outermost occupied shell (group A # between 1 and 8) - Octet rule o Main group elements react so that they attain octet with filled s and p subshells 3.4 Describe molecular shapes using VSEPR theory. 3.5 Explain chemical properties of elements based on the periodic table. 4.0 Ionic and Molecular Compounds - Picture 4.1 Explain the nature of ionic and covalent bonding. - Picture - Secondary bonding o Ionic dipole forces  Two atoms that have equal and opposite electrical charges  Picture o London forces  Between nonpolar molecules  Temporary dipoles form when more electrons happen to be on one side of the molecule than the other  Picture o Hydrogen bonding  Molecules containing hydrogen atoms that are covalently bonded to another highly electronegative atom (N, O, F)

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4.2 Explain the formation of ionic and molecular compounds. 4.3 Describe the general properties of ionic and molecular compounds. 4.4 Explain the bonding in metallic materials

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Nomenclature: o 1- mono o 2-di o 3-tri o 4-tetra o 5-penta o 6-hexa o 7-hepta o 8-octa o 9-nona o 10-deca Polyatomic Ions (13 ions) Oxyacid The number of valence electrons can be determined by the column number in the periodic table Mole Calculations

m (mass – g) = n (moles – mol) x M (molar mass – g/mol) o particles = n (moles – mol) x NA (avogadros number 6.02x1023 particles/mol) o C (molarity/molar concentration- mol/L or M) = n ((moles – mol) / V (volume – L) o P (pressure – kPa or atm) x V (volume – L) = n (moles – mol) x R x T (temperature – K)  R = 8.314 kPa L/mol K  R = 0.0821 atm L/mol K  SATP: 25ºC, 100kPa  STP: 0ºC, 101.3kPa  1 atm = 760mm of mercury (height) Solubility rules (solubility chart needed) o Li>K>Ba>Ca>Na>Hwater>Mg>Al>Zn>Cr>F e>Cd>Co>Ni>Sn>Pb>Hacid>Cu>Ag>Au Ionic equations o Balanced equation: Ca (s) + 2AgNO3 (aq) -> 2Ag (s) + Ca(NO3)2 (aq) o Ionic equation: Ca (s) + 2Ag+(aq) + continued on page Endothermic reaction absorbs energy from the surroundings Exothermic reaction release energy from the surroundings Calorimetry o q (thermal energy) = m (mass - g) x c (temperature - ºC) x Delta T (specific heat capacity – J/g ºC) Bond Dissociation Energy o Delta H equation continues Rates of Reactions o Factors affecting rates of reactions  Temperature – higher temp is faster  Concentration – higher concentration is faster  Nature of reactants  Surface area – higher surface area is faster  Catalyst – speeds up a reaction o Measuring reaction rates  Colour change  pH  pressure  volume of gas  mass Equilibrium o aA + bB = dD + eE o K = (Dd x Ee) / (Aa x Bb) -> products over reactants (exponents are the coefficients) o K >1 – lots of products o K = 1 – products = reactants o K< 1 – lots of reactants pH and pOH o pH = -log (H3O+) o pOH = -log (OH-) o pH + pOH = 14 o

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Intermolecular forces

Hydrogen bonding > dipole-dipole forces > London dispersions o Intramolecular force – force within a molecule o Intermolecular force – forces between molecules o Dipole-dipole – between polar molecules o London – occur between non-polar Atomic structure o Principle quantum number (n)  n = 1, n = 2, n = 3 o Secondary quantum number (l)  l = 0 is a s orbital  l = 1 is a p orbital  l = 2 is a d orbital  l = 3 is a f orbital o Magnetic quantum number (ml)  s orbital - ml is 0  p orbital - ml is -1, 0, +1  d orbital - ml is -2, -1, 0, +1, +2  f orbital - ml is -3, -2, -1, 0, +1, +2, +3 o Spin quantum number  + ½ or – ½ o

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Note: no two electrons in the same atom can have the same four quantum numbers

Step-by-step -

Wavelength using Balmer-Rydberg equation Atomic weigh...