Studied by 27 people
Briefly summarize the theory of the atom as stated by: Dalton, Thomson, Rutherford, Chadwick, and Bohr.
Dalton- All matter is made up of tiny, indivisible particles called atoms, all atoms of an element are identical, and atoms of different elements are different. Atoms are rearranged to form new substances in chemical reactions, but they are never created or destroyed.
Thomson- “plum pudding model”, used a cathode ray tube to discover the electron.
Rutherford- Aimed tiny Alpha particles at a thin sheet of gold foil. The positive particles are protons, the nucleus is surrounded mostly by space and electrons. Predicted the existence of the neutron. Nuclear model- an atom contains a small, dense positive central nucleus.
Chadwick- He confirmed with experiments that nuclei contain neutral particles as well as protons. These neutral particles are neutrons.
Bohr- Observed that hydrogen emitted light when excited by energy. Made the planetary model: electrons orbit the nucleus in definite energy levels (shells) and can jump between energy levels. Each energy level holds a specific number of electrons. The outer orbit of electrons is called the valence shell.
What is the difference between theoretical and empirical knowledge?
Theoretical knowledge is the knowledge that explains scientific observation, empirical knowledge is knowledge that comes from investigation and observation.
What is the relationship between the number of proton, neutrons and electrons in an atom?
Proton is the atomic number, electron has the same number as protons, and neutrons is mass- atomic #.
The variable “Z” is given to the atomic number of an element while the variable “A” is assigned to the atomic mass of an element. Two atoms respectively have a Z=12, A=26 and Z=14, A=26. Can these two atoms be classified as isotopes of the same element? Give reasons for your answer.
No, because the atomic number is different, and the atomic masses are the same. For these two atoms to be isotopes, the atomic number has to be the same, and the atomic masses are supposed to be different.
Natural magnesium exists as magnesium-24 with an abundance of 78.7%, magnesium-25 with an abundance of 10.1% and magnesium-26 with an abundance of 11.2%. Calculate the atomic mass that would be reported on the periodic table for magnesium using this data.
Define: ionization energy, electron affinity, electronegativity, atomic radius, effective nuclear charge (ENC), and shielding.
Ionization energy- Ionization energy is the energy required to remove an electron from an atom or ion. Must overcome attraction between the negative electron and positive nucleus. 2nd Ionization energy- The energy required to remove a second energy from a cation. More than 1st ionization since the nucleus is more positive after the first electron is lost. Group 1- loses its first valence electron easily, with low first ionization energy, and second ionization energy is very high since the new valence is full. Group 2- easily loses 2 electrons, low first and second ionization energies, and high third ionization energy because the new valence is full. Metals lose electrons!!!! Once an element has a full valence, the energy to break that stability is very high! Decreases as you move down a group, due to increased electron shielding and decreased ENC. Increases as you move across a period, due to an increase in ENC and unchanged shielding.
Electron affinity- the energy given off when an electron is added to an atom to make an anion (non-metal). Elements with high electron affinities form negative ions in ionic compounds. Elements with low electron affinities form positive ions in ionic compounds. Decreases as you move down a group, due to increased electron shielding, and decreased ENC. Increases as you move across a period, due to increased ENC, unchanged shielding (easy to gain e-).
Electronegativity- a measure of the ability of an atom in a compound to attract electrons in a bond higher electronegativity means a stronger pull on electrons Ranges from 0 to 4 (value written on periodic table below the atomic number for each element). decreases as you move down a group due to a decrease in ENC, lots of shielding increases as you move across a period increasing atomic number, so an increase in ENC, unchanged shielding. Metals have a low electronegativity, Non-metals have a high electronegativity.
Atomic radius- The distance from the center of an atom to the outermost electrons. Increases as you move down a group and decreases across a period.
Ionic radius- Cations are smaller than the atom they come from, ENC increases, and shielding decreases. Anions are larger than the atom they come from, ENC decreases, shielding unchanged, and repulsion increases. Ionic radius increases as you move down a group, due to decreased ENC as shielding is unchanged. Ionic radius decreases due to increased ENC as you move across a period.
Effective nuclear charge (ENC)- Charge felt by the valence electrons after you have considered the number of shielding electrons.
Electron shielding- The inner electrons protect the outer electrons from the attractive force of the nucleus.
Rank each of the following in order of increasing atomic size, give reasons for your ranking
a) Mg, S , Cl b) Al, B, In c) Ne, Ar, Xe d) Rb, Xe, Te
A) 1. Mg 2. S 3. Cl
B) 1. B 2. Al 3. In
C) 1. Ne 2. Ar 3. Xe
D) 1. Rb 2. Te 3. Xe
Rank each of the following in order of decreasing ionization energy, give reasons for your ranking
a) Na, Li, Cs b) S, Cl, Br c) Cl, Ar, K d) Ga, Ge, Se
A) 1. Li 2. Na 3. Cs
B) 1. S 2. Cl 3. Br
C) 1. Cl 2. Ar 3. K
D) 1. Se 2. Ge 3. Ga
Which element in the pair will have the lower electron affinity give reasons for your ranking
a) K or Ca b)O or Li c) S or Se d) Cs or Fr
A) K B) Li C) Se D) Cs
Calculate the difference in electronegativity for each bond and indicate the type of bond (non-polar covalent, polar covalent, ionic)
a) Zn-O b) Mg-I c) Co-Cl d) N-O
A) 1.9 Polar B) 1.3 Polar C) 1.2 Polar D) 0.5 Polar
Intermolecular forces exist BETWEEN two molecules. a) Explain what type of molecular polarity would result in London dispersion forces, dipole-dipole forces, or hydrogen bonding forces. b) Rank the strength of the intermolecular forces beginning with the strongest and ending with the weakest. c) Which force would increase the boiling point of a molecule? Why?
A) LDF- non-polar, dipole-dipole and hydrogen bond- polar.
B) Hydrogen bond, dipole-dipole, LDF
C) Hydrogen bonding because they have strong boiling points.
For each molecule below state: Central atom, total valence electron count, bond polarity, draw the lewis structure, state the molecular polarity and intermolecular forces present for that molecule.
a) SiO2
b) PCl5
c) H2O
d) NCl3
e) CF4
A) Si is a central atom, 16 e-, EN= 1.7 polar bond, non-polar molecule because symmetrical, and LDF, and dipole-dipole forces are present.
B) P is CA, 40 e-, EN= 0.9 so polar bond, non-polar because symmetrical, LDF.
C) O is CA, 8 e-, EN=1.4 polar bond, polar because asymmetrical, LDF and H-bond.
D) N is CA, 26 e-, EN=0 non-polar, non-polar because it’s a non-polar bond, LDF
E) C=CA, 32e-, EN= 1.8 polar, non polar because symmetrical, LDF.
Write the formula for the following compounds:
a) Silver chloride b) Phosphorus pentachloride c) Titanium (IV) oxide d) Calcium oxide e) Manganese (II) phosphide f) carbon tetrahydride g) Mercury (II) fluoride h) Iron (II) sulfide i) Dinitrogen trioxide
A) AgCl B) PCl5 C) Ti2O4 D) CaO E) Mn3P2
F) CH4 G) HgF2 H) FeS I) N2O3
Write the name of each compound, remember to watch out for the various types of compounds: ionic, molecular, hydrates, and acids as they are named differently.
a) HIO2 (aq) b) CsF c) NaOH (aq) d) K2Cr2O7 e) KClO4 f) N2Cl2 g) Al2(SO3)3 h) Fe(IO4)3 i) CuSO4∙5H2O j) Fe(IO4)3 k) MgCO3∙3H2O
A) Iodite B) Cesium Fluoride C) Sodium hydroxide
D) Potassium dichromate E) Potassium perchlorate
F) Dinitrogen dichloride G) Aluminum sulfite H) Iron periodate I) Copper sulfate pentahydrate J) Iron periodate K) Magnesium carbonate trihydrate
State the Law of Conservation of Mass. How does this relate to the reason why we balance chemical reactions?
Matter can’t be created or destroyed they can only be rearranged. We balance equations so that all of the products are all equal to the reactants.
Write a balance chemical equation for the reaction between the following elements/compounds and state the reaction type. Be sure to write the states of matter and consult the activity series/solubility table to ensure the reaction occurs.
a) Barium chloride reacts with Sodium carbonate
b) Iron reacts with copper (II) sulfate
c) ethyne (C2H2) reacts with oxygen gas. (do it for an oxygen rich environment and then an oxygen poor environment)
d) phosphorus pentachloride decomposes into its elements
e) Calcium reacts with chlorine gas
A) BaCl2 (s) + Na2CO3 (aq) → BaCO3 (s) + 2NaCl (s)
B) Fe (s) + CuSO4 (aq) → FeSO4 (aq) + Cu (s)
C) C2H2 (aq) + 2O2 (g) → CO2 (g)+ CO (g) + H2O (l) + energy + soot
or 2C2H2 (aq) + 5O2 (g) → 4CO2 (g) + 2H2O (l) + energy
D) 2PCl5 (aq) → 2P (s) + 5Cl2 (g)
E) Ca (s) + Cl2 (g)→ CaCl2 (aq)
What 3 scenarios will result in products being formed in a double displacement reaction?
Precipitation reaction, reactions that produce water, and reactions that produce a gas.
What scenario will result in no reaction for a double displacement reaction? Single displacement reaction?
When both products from a double displacement reaction are aqueous the reaction does not occur. For single displacement if the higher metal on the activity series is in the compound the reaction does not occur.
Find the molar mass of the following compounds:
a) LiF
b) Al2O3
c)Fe(IO4)3
d) CuSO4∙5H2O
A) 25.94 g/mol
B) 101.96 g/mol
C) 628.55 g/mol
D) 249.68 g/mol
How many moles are in a 78.1g sample of PH3
2.30 mol are in a 78.1g sample of PH3
How many atoms of silver are present in a 12.2g necklace?
6.81 × 10²² atoms of silver are present in a 12.2g necklace.
It was found that a sample of H2O contained 3.49 X 1023 molecules of water by calculation. How many atoms of H are present in this sample?
6.98 × 1023 atoms of Hydrogen in this sample.
In a reaction between sodium chloride and silver nitrate,
a) How many moles of NaCl would need to react with 1.4 mol of silver nitrate?
b) What mass of silver nitrate reacts with 29.2g of sodium chloride?
A) 1.4 mols of NaCl would need to react with silver nitrate.
B) 84.9g of silver nitrate reacts with 29.2g of sodium chloride.
Determine the percentage composition of CCl2F2
C= 9.9% Cl= 58.6% F= 31.5%
Researchers have isolated a compound that contains only nitrogen and oxygen. They found that a 4.60g sample of this compound contains 1.40g of nitrogen and the remainder is oxygen. Find the percentage composition for this compound.
N= 30.4% O= 69.6%
Progesterone, a hormone in the human body, is made up of 80.2% C, 10.18% O and 9.62% H. Determine the empirical formula for progesterone.
C21O2H30
An unknown compound that contains 64.3% carbon, 7.2% hydrogen and the remainder oxygen was found to have a molar mass of 168.21g/mol by lab analysis. Determine the empirical and molecular formula of this unknown compound.
C3H4O and C9H12O3.
Powdered zinc and sulfur react in an extremely rapid, exothermic reaction to produce zinc sulfide. A 6.00g sample of zinc is allowed to react with 3.35g of sulfur.
a) Which reactant was the excess reagent?
b)Which reactant is the limiting reagent? How do you know it is the limiting reagent?
c)Determine the theoretical yield of zinc sulfide
d) What % yield of ZnS if 6.63g of ZnS was obtained after processing.
e) What factors cause actual yield to be less than theoretical yield?
f) What may have occurred if your % yield is GREATER THAN 100%?
A) Zinc
B) Sulfur, I know that this is the limiting reagent because the limiting reagent is fully consumed so it is the substance that has less mass.
C) 8.94 g is the theoretical yield of zinc sulfide.
D) 74%
E) The nature of the reaction, impurities, competing side reactions, and the experimental procedure.
F) Other reactions could have formed and created more product, making the percentage yield greater than 100%.
What factors affect the rate of dissolving of a solid in a liquid?
Temperature and pressure can affect solubility.
Which of the following would you expect to be soluble in water?
a) KCl b) CCl4 c) Na2SO4 d) C4H10
KCl and Na2SO4 because they are ionic compounds.
0.25g of lead were found in a 4000.0L water sample. What concentration of lead, in ppm, is present in the sample?
C=0.063 ppm
A 250mL sample of waste sewer water is found to contain 12 ppb of covid-19 viral protein. Determine the mass of covid-19 protein found in the sample.
The mass of the covid-19 viral protein is 3.0g.
Calculate the mass of solute that is needed to prepare each solution below:
a) 250mL of 0.250mol/L calcium acetate
b) 1.8L of 0.35mol/L ammonium sulfate
A) 9.9 g
B) 83g
Calculate the molar concentration of a 70.0mL solution produced by diluting 20mL of 6.0mol/L hydrochloric acid stock solution.
2.0 mol/L
Calculate the volume of 12.0mol/L stock solution required to produce a dilute solution of ammonia with a volume of 2.50L and concentration 1.44 mol/L
0.3 L
Hydrogen sulfide gas can be prepared by the reaction of sulfuric acid with sodium sulfide. Write the net ionic equation for this reaction.
2H + (aq) + S 2- (aq) → H2S (g)
50mL of 0.200mol/L Barium nitrate is mixed with 200mL of 0.180mol/L potassium sulfate. What is the theoretical yield of the barium sulfate solid produced by the reaction?
2g
A double displacement reaction occurs in an aqueous solution when magnesium phosphate reacts with lead (II) nitrate. If 20.0 mL of 0.750 mol/L magnesium phosphate reacts, what is the maximum mass of the precipitate that can be formed?
12.2g
Phosphoric acid is reacted with sodium hydroxide. Given the laboratory data below, calculate the unknown concentration of the acid:
Trial | 1 | 2 | 3 | 4 |
Vol of acid (unknown conc.) | 10.0 mL | 10.0 mL | 10.0mL | 10.0mL |
Conc. Of base (titrant) | 1.25 mol/L | 1.25 mol/L | 1.25mol/L | 1.25mol/L |
Final burette reading base | 13.3 mL | 25.0 mL | 36.8 mL | 48.4 mL |
Initial burette reading base | 0.4 mL | 13.3 mL | 25.0 mL | 36.8 mL |
Vol. Base added | ||||
Colour at endpoint | Deep red | Pale pink | Pale pink | Pale pink |
0.488 mol/ L
A pressure gauge reads a pressure of 310 atm what is the pressure in kPa?
310 KPA
A temperature is read to be -78 ⁰C what is this temperature in Kelvin?
195 K
What volume would 1.0mol if O2 occupy at STP? SATP?
V= 22 L V= 25 L
A sample of oxygen gas occupies a volume of 10.0L at 546K. At what temperature in ·C would the gas occupy at a volume of 5.0L? Assume P and n are constant
T2= 0 ·C
A sample of nitrogen gas occupies 11.20L at 0·C and 101.3kPa. How many moles of nitrogen are there in this sample?
0.500 mol
A 250.0mL balloon full of pure helium at 101.3 kPa is subjected to a pressure of 125.0kPa at a constant temperature. What is the final volume of the balloon? Assume T and n are constant.
V2= 0.203 L
The head of a match contains approximately 0.75g of diphosphorus trisulfide. When the match is struck on a rough surface, it explodes into flames producing diphosphorus pentaoxide and sulfur dioxide. What volume of sulfur dioxide will be produced if the temperature is 26.5· C and the pressure is 102.8kPa?`
V= 0.34 L
What is the partial pressure of hydrogen gas if it is part of a mixture of Oxygen gas with a partial pressure of 10 kPa and nitrogen gas with a partial pressure of 35 kPa if the overall pressure exerted by the mixture is 86 kPa?
Partial pressure is 41 kPa.
27.5 mL of Hydrogen gas is collected using downward displacement of water during an experiment between magnesium metal and hydrochloric acid. If the experiment is conducted at a pressure of 100.50 kPa and a temperature of 20.0 ⁰C , determine
a) the partial pressure of the hydrogen gas
b) the number of moles of hydrogen gas that was collected.
A) P= 98.16 kPa H2
B) n= 0.001 mol H2
A 550 mL sample of methane, CH4 (g) is collected over water by downward displacement. The temperature is 25 ⁰C and the atmospheric pressure is 100.7 kPa. What is the partial pressure of the methane? What mass of methane was produced?
PCH4= 97.53 kPa
mCH4= 0.35g
Ground state
This is when electrons in an atom have the lowest amount of energy possible for the location that they are at in the element
Excited state
When energy is added in form of heat or electricity to an atom, electrons absorb energy which allow them to “jump” to a higher energy level.
emission spectrum
each element has a unique characteristic pattern of electron movement
Isoelectronic
isoelectronic is when an element has the same number of electrons as a noble gas on the p.t.
Isotope
atoms of the same element that have different number of neutrons, protons and electrons don’t change.
Isotopic abundance
the percentage of a given isotope in a sample of the element.
Radioisotopes
Some isotopes are stable but radioisotopes aren’t because large numbers of neutrons destabilizes the nucleus.
Types of radiation
alpha particles- same structure as helium nucleus w a 2+ charge, blocked by paper
beta particles- negatively charged electron, blocked by aluminum.
gamma rays- have no mass and travel at light speed, blocked by lead
alkali metals
soft silver solids, reacts violently in water to make a base. hydrogen is not apart of alkali metals
alkaline earth metals
solid, reacts w air to form oxide coatings, all except Be react with hydrogen to form hydrides.
halogens
can be any state, dull and don’t conduct electricity, extremely reactive
noble gases
gases, low melting and boiling point, unreactive (inert)
polyatomic ion
a group of nonmetal atoms that have an overall ionic charge, names end in -ate or -ite.
hydrate
An ionic compound that includes water as part of its crystal structure.
difference between ionic and molecular compounds.
An ionic compound is a compound of a metal and a non-metal and the elements lose or gain electrons. Ionic compounds are solids, with high melting and boiling points, and dissolve in water easily. Molecular compounds are made of two non-metals, they share electrons. Molecular compounds can be solid liquid or gas, low melting and boiling points, don’t dissolve in water easily.
acid
an acid is a substance containing hydrogen and is in the aqueous state (aq) which means its dissolved in water. Two types: binary and oxyacid, binary is a hydrogen and non-metal, or hydrogen and a polyatomic ion that doesn’t contain oxygen. Oxyacid- hydrogen and polyatomic ion with oxygen.
Unique properties of water
high melting and boiling points- allows water to be a liquid at room temp
expanding when cooling from 4-0 oC- expansion causes ice to float, causing water to freeze from the top allowing life to live below it.
high surface tension- pulls water in round droplets allows small insects to walk on water.
ability to exchange thermal energy with little temp change- allows water to absorb or release a lot of thermal energy for a small increase or decrease in temp.
inability to mix with non-polar molecules- enables organisms to retain water because of a waterproof coating, allows organisms to store non-polar substances.
chemical change
the transformation of one or more substances into different substances with different properties.
chemical reactions
energy is absorbed or released, occur at different rates (slow or fast) catalyst speeds up reactions
complete combustion
lots of oxugen
fuel + oxygen → CO2 + H2O + energy
incomplete combustion
limited oxygen
fuel + oxygen → C + CO + CO2 + H2O + energy + soot
solution
homogeneous mixture of a solute dissolved in a solvent.
Oxides
a compound made up of any element and oxygen. metal oxide are called basic oxides. non-metal oxides are called acidic oxides.
decomposition reactions
Δ= heat
Hv= electricity
Formula of catalyst like MnO2 or the word catalyst.
double displacement
in the case that the double displacement reaction produces a gas if it is H2CO3 it is broken up into H2O and CO2.
gas tests
oxygen- glowing splint ignites
hydrogen- flaming splint pops
carbon dioxide- flaming splint extinguishes
water- cobalt chloride paper changes from blue to pink
decomposition involving compounds
chlorate breaks down into a metal chloride and oxygen gas, carbonate breaks down into a metal oxide and a carbon dioxide.
what are the components of a mixture
solute- the substance that has dissolved (lesser quantity)
solvent- the substance that does the dissolving (greater quantity)
types of solutions
alloy- solution of two or more metals
amalgam- an alloy of mercury
aqueous solution- solution with water as a solvent
what is meant by like dissolves like
polar dissolves polar, non-polar dissolves non-polar.
surfacants
compounds with a hydrophilic end attracted to water, and a hydrophobic end that is repelled by water. these compounds reduce the surface tension of a solvent in order to make it more soluble.
define under saturated, saturated and super saturated
undersaturated- a solution containing less solute than it can typically hold at a given temp and pressure
saturated- a solution can’t dissolve any more solute at the given temp and pressure.
super saturated- a solution contains more solute than it can typically hold at a given temp and pressure
how to read the solubility table
saturated- point on the curve
undersaturated- point under curve
supersaturated- point above curve
define crystallization
crystallization- process used to recover certain solutes from a solution.
precipitation- type of reaction used to identify unknown solutes in solution.
solubility of solids liquids and gases in water
solids- temp goes up, solubility goes up.
gases- temp goes up, solubility goes down.
liquids- hard to predict
concentration
quantity of solute per unit volume of solution.
how to write net ionic equations
total ionic equation- ionic compounds dissociated as ions, solids liquids and gases don’t dissociate.
net ionic equation- only shows the ions involved in the reaction
spectator ions- ions equal in quantity and have the same states of matter on both sides of the reaction.
stoichiometry of solutions
start with balanced equations, find moles using c and v, mole ratio, moles of required substance l, find c, v or m of required substance.
properties of acids and bases
acids- sour, no special feel, neutralizes bases, releases hydrogen gas and carbon dioxide, pH less than 7.
base- bitter, slippery, neutralizes acids, makes carbonates, pH more than 7.
naming acids and bases
binary acid= hydro + _____
oxyacid= change ending of polyatomic ion from -ate to -ic and -ite to -ous.
base- metal name + hydroxide
arrhenius theory
electrolytes conduct electricity because they dissociate into ions, since acids and based conduct electricity they must be made up of ions. molecular compounds ionize, ionic compounds dissociate.
define titration, titrant, burette, equivalance point, endpoint, and indicator.
Titration- A procedure used to determine the concentration of a solution using a standardized solution.
Titrant- The solution in the burette during a titration.
Burette- A calibrated tube used to deliver variable known volumes of a liquid during a titration.
Equivalence point- The point in a titration when neutralization is complete.
Endpoint- The point during a titration when a sudden change in an observable property of the solution occurs; usually a change in the colour of an acid-base indicator or a significant change in pH.
Indicator- The indicator matches the equivalence point, so that the colour change when a small excess of titrant is added.
gas stoichiometry
write balanced chemical equation, find moles of product whose info is given use whichever moles equation you can use. remember if you need to find moles of both reactants you must determine limiting reagent, use mole ratio to find moles of required substance. use PV=nRT to find quantity of interest, if gas is collected over water remember to use daltons law first to find P gas made first.
Note
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