Chemistry Flashcards
naming the acid when it has only an element following the H
use the prefix hydro-, followed by the element’s root name and an-ic ending
If the acid has an-ate polyatomic ion after the H
that makes it an-ic acid. H2SO4 is sulfuric acid
When the acid has an-ite polyatomic ion after the H
makes it an-ous acid
Atoms make bonds because
of their closed shell electron configurations
why can’t nobel gases accept another electron into the shell
Because the valence shell of a noble gas is completely full
the loss of an electron from a noble gas is unfavorable because
The nucleus is positively charged and pulls on the electrons
Noble gases are
inert/unreactive
Any element other than a noble gas has
an open shell configuration, which is unstable relative to the configuration of a noble gas.
Non-noble atoms react to form bonds in an attempt to
achieve a closed shell electron configuration
Even though the reaction may appear to be favorable because of its production of a closed shell species
there is a way to have both F atoms achieve a noble gas configuration.
covalent bonds also occur between
diatomics
sharing of electrons is called
- covalent bond
- 2 nonmetals
When a highly electronegative atom and an electropositive one are bonded together
an electron is transferred from the electropositive atom to the electronegative atom to form a cation and an anion, respectively.
how does distance affect attraction between the two ions
At large distances, there is a negligible energy of attraction between the two ions, but as they are brought closer together, they are attracted to one another.
how does distance affect ion attraction?
Ions are actually repelled at small distances. To explain this observation, remember that the ions’ nuclei are both positively charged. When the nuclei approach each other, they repel strongly–accounting for the steep rise in potential as the ions get closer than the bond length.
Ionic compounds are a part of
a crystal lattice–a three dimensional regular array of cations and anions.
why do Ionic compounds form lattices?
due to the contributing coulombic attractions of having each cation surrounded by several anions and each anion surrounded by several anions.
why are covalent bonds stable?
- due to the build-up of electron density between the nuclei.
- By sharing electron pairs nuclei can achieve octets of electrons in their valence shells, which leads to greater stability.
why does a Lewis structure only counts valence electrons
because these are the only ones involved in bonding. (To calculate the number of valence electrons, write out the electron configuration of the atom and count up the number of electrons in the highest principle quantum number.)
We can create bonds by having
two atoms come together to share an electron pair.
how is a bonding pair of electrons is distinguished from a non-bonding pair
by using a line between the two atoms to represent a bond (A lone pair is what we call two non-bonding electrons localized on a particular atom.)
In atoms, electrons reside in
orbitals of differing energy levels such as 1s, 2s, 3d, etc.
orbitals represent
the probability distribution for finding an electron anywhere around the atom.
Molecular orbital theory
posits the notion that electrons in molecules likewise exist in different orbitals that give the probability of finding the electron at particular points around the molecule.
bond order
- the number of bonds between atoms in a molecule.
- The bond order is the difference in the number of electron pairs occupying an antibonding and a bonding molecular orbital.
Steps for Basic Stoichiometry Calculations
4 Steps:
- Balance the equation
- Convert units of given substance to moles
- Find moles of wanted substance using mole ratio
- Convert moles of wanted substance to desired units
Moles =
grams/formula mass
Converting between gas volume and moles
Number of moles = PV/RT
At STP, a mole of gas will always occupy
22.4 L of volume
Avogadro’s Number
- provides the conversion factor for moving from number of particles to moles.
- There are 6.02×1023 formula units of particles in every mole of substance
Limiting Reactant Problems
- Whichever reactant that limits the production of product is the limiting reactant
- Start by converting to molesO2(g)
- Figure out how many grams of C(s) react. Mole ratio tells us that an equal number of moles will react.
- Subtract to figure out how much C(s) is left over.
- limiting reagent limits or determines the amount of product that can be formed.
Percent Yield
- (actual yield/theoretical yield)*100
- ## How many grams of CaO should be produced? First verify the equation is balanced. Now convert to moles, based on the amount ofCaCO3 present, then to grams
Energy Changes
- ## Reactions that release energy in the form of heat are called exothermic reactions. Conversely, reactions requiring heat energy are known as endothermic reactions.
Based on the following balanced equation, how many kcal of energy are needed to decompose 300 grams of CaCO3(s) ?
CaCO3(s) + 176 kJ→CaO(s) + CO2(g)
- Convert to moles.
- 300g/100g ×1 mole CaCO3 = 3 moles CaCO3
- Now do the mole ratio; 176 kJ are needed for every mole of CaCO3 .
- 3 molCaCO3/1 mol CaCO3 ×176 kJ = 528 kJ
- put the answer in terms of kcal. There are 4.18 kJ in every kcal. Use the necessary conversion factor.
- ## 528kj x 1.00kcal / 4.18kj = 126 kcal
Enthalpy
how much heat a substance has at a given temperature and pressure, and is symbolized by the symbol H
heat of reaction
- The change in enthalpy for a reaction
- has symbol δH
enthalpy determines
- exothermic or endothermic
- δH is negative for all exothermic reactions
- positive for all endothermic reactions
Heat of formation equation
δH = δH f (of all products) - δH f (of all reactants)
Pressure=
Force/Area
Boyle’s Law
- Boyle’s law: P 1 V 1 = P 2 V 2
- downward curve on volume v pressure graph
- The most important thing to remember about Boyle’s Law is that it only holds when the temperature and amount of gas are constant.
- A state of constant temperature is often referred to as isothermal conditions.
- temp doesn’t have to be in kelvin
Isothermal conditions
A state of constant temperature
The Manometer
- There are two ends to Boyle’s manometer. One end is open to the atmosphere. The other end is sealed, but contains gas at atmospheric pressure. Since the pressure on both ends of the tube is the same, the level of mercury is also the same.
Manometer J-shaped, closed on curved end with V-50ml and open on straight end (Patm)
- The pressure of the gas before mercury is added is equal to the atmospheric pressure, 760 mm Hg
- After added mercury, the volume of the gas, V 2 , drops to 50 mL
- P 2=P 1 V 1/V 2=
(100 mL)(760 mm Hg)/(50 mL)=1520 mm Hg
Charle’s Law
- V1/T1 = V2/T2
- If temperature is measured on a Celsius scale, T can be negative. The standard absolute scale is the Kelvin (K) scale. The temperature in Kelvin can be calculated via T k = T C + 273.15 .
- straight line up on volume v temp graph
Density=
PM/RT (where M is molar mass)
Partial Pressure=
- P tot = P A + P B + P C + …
- moles g A/total mols = PP gas A/total pressure
Each individual gas obeys the ideal gas law so
- we can rearrange PV= nRT to find pressure:
- P1=n1RT/V
Arrhenius
defined acids to be proton (H+) donors and bases to be hydroxide ion (OH-) donors in aqueous solution.
Bronsted-Lowry
- describing acids as proton donors and bases as proton acceptors.
- The Bronsted-Lowry model implies that there is a relationship between acids and bases (acids transfer protons to bases) and allows defines conjugate acids and conjugate bases
- HA (acid) + B (base)= A-(CB) + BH+(CA)
Lewis model
proposes that an acid is an electron pair acceptor while a base is an electron pair donor.
pH =
- log [H+]
pOH
the negative common logarithm of the concentration of OH-
pH + pOH =
14
pH, pOH, [H+], and [OH-] at extreme acid
- pH=14
- pOH=0
- [H+]=10^-14
- [OH-]=10^0=1
When a strong acid or a strong base is added to water
it nearly completely dissociates into its ion constituents because it has a pKa or pK b less than zero. For example, a solution of H2SO4 in water contains mostly H+ and SO4 2
The concentration of acid equals
the concentration of H+
Common strong acids that should be memorized
- Hydrochloric HCl
- Hydrobromic HBr
- Hydroiodic HI
- Nitric HNO3
- Sulfuric H2SO4 (only strong first time)
- Perchloric (HClO4)
Strong bases include
Group I hydroxides (LiOH, NaOH, KOH, etc.) and Group II hydroxides except for Be(OH)2 and Ba(OH)2.
why is calculating the pH of weak acid and weak base solutions is much more complicated than strong?
weak acids and bases do not completely dissociate in aqueous solution but are in equilibrium with their dissociated forms.
calculate the pH of a 0.10 M solution of acetic acid in water
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Polyprotic Acids
acids that can donate more than one proton per molecule.
Two key features of polyprotic acids
- they lose their protons in a stepwise manner
- each proton is characterized by a different pK a.
factors contributing to the pK a of each acidic proton in a polyprotic species
The factors contributing to the pK a of each acidic proton in a polyprotic species are the same factors that determine the relative acidity of monoprotic acids–the dominant factor is strength of the acid-H bond.
Buffers
- A buffer is simply a mixture of a weak acid and its conjugate base or a weak base and its conjugate acid
- Buffers work by reacting with any added acid or base to control the pH.
- A buffer works by replacing a strong acid or base with a weak one
Calculating the pH of Buffered Solutions
Use the Henderson-Hasselbalch Equation
Henderson-Hasselbalch Equation
- pH=pKa+log(base/acid)
- pH= PKa + log (A-/HA+)
- POH=PKb + log (HA/A)
- Find pKa using the expression that’s similar to equilibrium
Titrations
- A titration curve is drawn by plotting data attained during a titration, titrant volume on the x-axis and pH on the y-axis.
- equivalence point at steepest part
Redox Reaction
The oxidation state of an atom in a covalent compound is an imaginary charge assigned to that atom if all the electrons in its bonds were completely given to the more electronegative atom in the bond.
redox rules
- Atoms in elemental form have oxidation states of zero.
- The charge on a monoatomic ion is equivalent to its charge.
- Hydrogen has an oxidation state of +1 when bonded to non-metals and -1 when bonded to metals.
- F, because it only forms one bond and is the most electronegative element, has an oxidation state of -1.
- O, unless bonded to F or itself, has an oxidation state of -2.
- The sum of all oxidation states in a compound must equal the total charge on the species. For hydrogen peroxide, we have seen that both O’s have -1 oxidation states and both H’s have +1 oxidation states, the sum of which is zero–the charge on hydrogen peroxide.
Balancing Redox Reactions
- Separate oxidation and reduction half-reactions:
- Balance all atoms except for hydrogen and oxygen in each half-reaction. In this example they are already balanced.
- Balance oxygen by adding H2O as needed:
- To balance hydrogen, add H+ as needed: (Note: You still do this if you are in basic solution, later on, you will add OH- to “neutralize” the acid.)
- Balance the charge of each reaction by adding electrons to side with the greater charge:
- Multiply each half-reaction by the least integer factor that equalizes the number of electrons in each half-reaction. Then, add the half-reactions to obtain the overall balanced reaction in acidic solution:
- If your redox reaction is in acidic solution, the above reaction is properly balanced. However, if the reaction you wish to balance is in basic solution, you need to add these three steps:
- If the redox reaction is one in basic solution, then add OH- to both sides of the equation to “neutralize” each H+:
- “React” H+ and OH- to form H2O and eliminate water molecules on both sides of the equation
- Make sure that all atoms and charges are indeed balanced in your overall balanced equation for the redox reaction in basic solution
Galvanic Cells
- Galvanic cells harness the electrical energy available from the electron transfer in a redox reaction to perform useful electrical work. The key to gathering the electron flow is to separate the oxidation and reduction half-reactions, connecting them by a wire, so that the electrons must flow through that wire.
two typical setups for galvanic cells
- The left hand cell diagram shows and oxidation and a reduction half-reaction joined by both a wire and a porous disk
- The right hand cell diagram shows the same cell substituting a salt bridge for the porous disk.
Why is the salt bridge/porous disk necessary?
The salt bridge or porous disk is necessary to maintain the charge neutrality of each half-cell by allowing the flow of ions with minimal mixing of the half-cell solutions. As electrons are transferred from the oxidation half-cell to the reduction half-cell, a negative charge builds in the reduction half-cell and a positive charge in the oxidation half-cell. That charge buildup would serve to oppose the current from anode to cathode– effectively stopping the electron flow
reduction/oxidation in relation to anode/cathode
- reduction takes place at the cathode and oxidation takes place at the anode.
- “The Red Cat ate An Ox”
anode/cathode marked
The anode, as the source of the negatively charged electrons is usually marked with a minus sign (-) and the cathode is marked with a plus sign (+).
Negatively charged electrons flow in a wire. Therefore
chemists indicate the direction of electron flow on cell diagrams and not the direction of current.
cell potential, Ecell of a galvanic cell is measured in
volts
how do we measure the Eo of any half- reaction.
by arbitrarily assigning a value of exactly zero to the potential of the standard hydrogen electrode
E Cell example
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Electrolytic cells
like galvanic cells, are composed of two half-cells–one is a reduction half-cell, the other is an oxidation half-cell.
the direction of electron flow in electrolytic cells
the direction of electron flow in electrolytic cells may be reversed from the direction of spontaneous electron flow in galvanic cells, the definition of both cathode and anode remain the same–reduction takes place at the cathode and oxidation occurs at the anode.
comparing a galvanic cell to its electrolytic counterpart
When comparing a galvanic cell to its electrolytic counterpart, as is done in , occurs on the right-hand half-cell. Because the directions of both half-reactions have been reversed, the sign, but not the magnitude, of the cell potential has been reversed.
The electrolytic cell reaction is not the only one occurring in the system-
- the battery is a spontaneous redox reaction.
- By Hess’s Law, we can sum the ΔG of the battery and the electrolytic cell to arrive at the ΔG for the overall process. As long as that ΔG for the overall reaction is negative, the system of the battery and the electrolytic cell will continue to function. The condition for ΔG being negative for the system (you should prove this for yourself) is that Ebattery is greater than - Ecell.
Hess’s Law
regardless of the multiple stages or steps of a reaction, the total enthalpy change for the reaction is the sum of all changes. This law is a manifestation that enthalpy is a state function.
Electrolysis of Water
- That spontaneous direction of reaction can be used to create water and electricity in a galvanic cell (as it does on the space shuttle). However, by using an electrolytic cell composed of water, two electrodes and an external source emf one can reverse the direction of the process and create hydrogen and oxygen from water and electricity.
- The reaction at the anode is the oxidation of water to O2 and acid while the cathode reduces water into H2 and hydroxide ion.
- requires input of energy because the products possess more chemical PE than H2)
The Electrolysis of Water is
The electrolysis of pure water is -1.23 V. To make the electrolysis of water occur, one must apply an external potential (usually from a battery of some sort) of greater than or equal to 1.23 V. .
Gibbs free energy
- Gibbs free energy
- DeltaH-TdeltaS
- -G=spontaneous forward reaction
- +G= spontaneous reverse reaction
- G=0=equalibrium
- DETERMINES SPONTANEITY
- According to Hess’s Law, only state functions (G, H, S) and not path functions (w, q, E) of a series of reactions may be summed to generate a new value for the overall reaction.
- You can add the ΔG or the ΔH of reactions using Hess’s Law
When the temperature of a 20g sample of water is increased from 10C to 30C, the heat transferred to the water is
- q=mcdeltaT
- c is about 1cal/g
- so 20g=20C
- (20g)(1cal)(20C)=400
An aqueous solution with pH5 at 25C has a hydroxide ion (OH-) concentration of
1x10^-9 molar
The volume of water vapor required to produce 44.8 liters of oxygen by the below reaction is
2H2O(g)–>2H2(g)+O2(g)
- 2 moles H2O to 1 mole O2
- volume of gas will also be 2:1
- so 89.6 liters H20 are required to produce 44.8 liters O2
When 190 grams of MgCl2 are dissolved in water and the resulting solution is 500 ml in volume, what is the molar concentration of MgCl2 in the solution?
- M=mol/l
- find mass of 1 mol of MgCl2
- determine moles of MgCl2 by dividing the mass of the sample by the molar mass of the compound (190/95.4=2 mol
- the mol/v= 2/.5=4M
a 600 ml container holds 2 moles of O2(g), 3 moles of H2(g), and 1 mole of He(g). Total pressure within the container is 760 torr. What is the partial pressure of O2?
- mol gas a/total mols=pp gas a/total pressure
- 2/6=x/760
- x=253
An ideal gas has a volume of 10 liters at 20C and a pressure of 750mmHg. Whats the volume of the same gas at STP?
PV=nRT
- initial PV/T=final PV/T
- (750)(10)/293=760V/273
- 10 X 750/760 X 273/293
When 3 moles of Fe2O3 are allowed to completely react with 56 grams of CO according to the above equation, approximately how many moles of iron, Fe, are produced?
Fe2O3(s)+3CO(g)–>2Fe(s)+3CO2(g)
- find the limiting reactant
- 56g CO=2 mol CO
- Fe2O3:CO=1:3
- CO is limiting
- CO to Fe is 3:2, so 2 moles produces aout 1.3 moles of Fe
organic compounds properties
- organic compounds are much more soluble in nonpolar solvents
- organic compounds don’t dissociate in solution
- poor conductors of electricity
isomers
same chemical makeup with identical constituent elements, but are arranged in a different geometrical arrangement,a and have different chemical properties
hydrocarbons
- the simplest organic compounds
- contain only carbon and hydrogen
- 3 categories:
1. alkanes- contain only single C-C bonds
2. Alkenes- contain double C-C bonds
3. Alkynes- contain triple C-C bonds
Alkanes
- saturated hydrocarbons because each carbon atom is bonded to as many other atoms as possible
- formula based on prefixes (CnH2n+2)
prefixes
- indicate the number of carbons in the hydrocarbon
- meth=1
- eth=2
- prop=3
- but=4
- pent=5
- hex=6
Alkenes
- unsaturated hydrocarbons because each carbon atom is not bonded to as many atoms as possible
- Formula (CnH2n)
suffixes
- indicate what types of C-C bonds are present
- ane=single bonds
- ene=at least 1 double bond
- yne=at least 1 triple bond
alkynes
- unsaturated hydrocarbons
- formula (CnH2n-n)
hydrocarbon rings
aromatic hydrocarbons, the simplest of which is benzene C6H6
Alcohol
- an O-H bonded to a carbon atom. Because of hydroxyl (OH-) group, alcohols are polar
- ol
- ex. Methanol
Halides
- Halogen bonded to a carbon atom
- can by named by suffixes: fluoride, chloride, bromide, iodide, or by prefixes: fluoro, chloro, broom,
- ex. chloromethane
