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Sharing of an electron pair (no transfer like in ionic)
- Usually between two non-metals
- Tend to obtain stable e.c.
Lewis Structures
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Remember to put formal charge (imaginary ) → try to have a formal charge of 0
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Formal charge = group number (1-18) - # of e in lone pairs – ½ (# of e in bond pairs)
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Truth #1: Elements in the third period and beyond can have more than 8 electrons per shell
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Boron can have only 6 and still be ok
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C, N, O, F are goodie-goods. Always follow the rules
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Resonance is when there is ambiguity in where the double bond is called Resonance
Bonding Continum
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Electron pairs not equally shared
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Use electronegativity values to predict whether is ionic or covalent
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If more than 1.7 in difference in e.n. then it is ionic
e.n. 0 ----- 1.7 ----- 3.3 → structural formula
Pure covalent ‘--- ionic---> *note that between 0-1.7 is polar covalent
e.g. HCl e.n.(H) = 2.1 e.n.(Cl) = 3.0 e.n.=0.9 Polar Covalent Bond
Polar Covalent Bonds
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Unequal sharing of electrons (part ionic, part covalent)
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One atom is partially negative and the other positive
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Electrons spend more time around atom with higher e.n.
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Polar Molecules – molecules that have a slightly positive and a slightly negative end.
- Consider bond polarity and molecular symmetry shape
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Use the +---------> to tell which is the polar side (arrow pointing toward negative end)
NOTE: breaking bonds require energy. Forming bonds releases energy
Intermolecular Forces
Strong bonds
- Ionic Crystal – ionic bonds as bonding force
- Covalent Crystals – covalent bonds to each other
Molecular Crystals – uses weak bonds have low melting and boiling points. Use covalent bonded atoms
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London Dispersion Forces : between non-polar molecules due to attraction from the nucleus and the electrons of a nearby molecule → momentary dipole created
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Dipole-Dipole: between polar molecules. Caused by a slight attraction on each ends of the molecules. Very weak though.
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Hydrogen Bonding: Occurs when H is on a molecule such as F, O, or N. Stronger than dipole—dipole. Unusually high boiling point and melting point. Hydrogens bond onto the negative poled molecule.
Valency
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The repulsion theory VSPER Is used to predict geometries of molecules
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The Main Postulate: structure around a given atom is determined principally by minimizing electron pair repulsions
- Gets as far away as possible (lone pairs have strongest repulsions)
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Activity Series – used for single displacements reactions
- Shows the reactivity strengths between elements (which would displace the other)
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Solubility – if in the solubility chart the substance has low solubility in it then a precipitate most likely has formed which means a reaction has probably occurred.
- Types of Reactions – combustion, single displacement, double displacement, neutralization, synthesis.
REDOX Reactions
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Oxidation – where the atom LOSES electrons
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Reduction – where the atom GAINS electrons
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VERY IMPORTANT** - the oxidizing agent GAINS electrons but the oxidant LOSES electrons
- the reducing agent LOSES electrons but the reductant GAINS electrons
Law of __________.
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Definite Proportions: A compound always contains same elements in the same ratios
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Multiple Proportions: Two elements combine in the same ratios always.
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Relative Atomic mass: Carbon 12.0 on the dot. One C-12 is given 12u
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Isotope Abundance: Percent of an isotope in an element sample
Balancing Nuclear Equations
-
Three types of radioactive particles:
-
Alpha [α]- helium particle He
-
Beta [β] e-
- Gamma – energy
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Protons He and Neutrons n can also be emitted during a nuclear reactions
-
Fission is breaking apart, and Fusion is coming together.
Unit 2: Liquids
Solutions
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Solvent is greater quantity and solute is less in quantity
-
Remember: Solubility of solid goes up when temperature goes up
-
Solubility of gasses goes down when temperature goes down
- Immiscible – don’t mix and miscible – mixes together
Concentration
- Measure of solute dissolved in the solvent
- Concentrations is measured in moles per litre [mol/L]
-
Formula is [ ] = C = n/v
-
Where [ ] is molar concentration of, C is concentrations, n is moles per v, volume
-
Remember conversions (for ppm questions), 1g/10mL (is the basis for converting)
How to Prepare Solutions
- Remember the moles are always in mol/L. Use this to then convert to the specified amount
-
Use formula CV = CV to determine proportional amounts
Reactions of Ions in Solutions
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Dissociations – is decomposing an ionic solid into its separate ions
-
Ionic Equation – when all the compounds are written as ions
-
Net Ionic Equations – when a balance chemical equation, write the ionic equation, the you take out the ions that didn’t react (remains aqueous) and you write the reacting ones in ionic equations
-
The ones that didn’t react are known as spectator ions
Acids and Bases
- The Browsted-Lowry Theory
-
Acids give up an hydrogen ions, or the proton donor
-
Bases on the other hand reacts with the hydrogen, or the proton acceptor
- Calculating pH and solutions
-
pH = -log[H] and 10= [H]
Unit 3: Gasses
Boyle’s Law
-
PV= PV where T is a constant. Use this formula when only P and V is given without mention of temperature change
Charles Law
-
is used when P is a constant. Use this formula when only V and T are given
Combined Gas Law
-
when all variables are in play.
-
Usually gives u one set of variables and gives u two from the other set
Molar Volume
- 1mol = 22.4L at STP (which is 273K and 103.5kPa)
Ideal Gas Law
-
PV = nRT → used usually when asking for moles or when moles are mentioned
-
Note that R = 8.314kPa/L•K
Partial Pressure
-
++…
-
In other words all the pressures of the gasses (the partial pressures) add up to the total pressure
-
Method: When given the compounds, use ideal gas law to calculate the pressure of each gas and then add them up to get the total pressure
Unit 4: Hydrocarbons
Organic Compounds
-
Are compounds that contain carbon (except CO, CO, and HCO)
-
The carbon containing compounds has to have C-C bonds
-
Carbonates and cyanides are not considered organic
- Characteristics of Organics
- Burnable
- Soft, and usually do not have a hard mineral structure
- Usually liquids or solids MP<300
Hydrocarbons
-
Contains only hydrogen and carbon (hence the name)
-
Homologous series – when one compound differs from a preceding one by a -CH group
-
Alkanes are saturated compounds meaning they have full single bonds
-
The general formula for alkanes is
Naming those Hydrocarbons
Naming Continued
- Cyclo – when a ring is formed
- Alkenes – when there is a double bond formed
- Alkynes – when there is a triple bond formed
Hydrocarbon Reactions
- Hydrocarbons are non-polar (use of London-dispersion forces)
-
Cracking – when the hydrocarbon breaks up into smaller alkanes
-
Reforming – smaller alkanes combine to larger alkane
-
Combustion – basically full combustion: alkane + → +
Heat Reactions
-
Exothermic – energy is released
-
Endothermic – energy is absorbed
-
Basic heat equation is q = mc where q = energy in J, m = mass of object, c = specific heat capacity and delta t = change in temperature. This is the general equation
-
There is also , generally use when combustion is mentioned or when dealing with calorimetry