Chemistry Exam 1 Study guide
∙ The first 30 elements***
∙ Chemistry is the study of matter.
∙ Matter = anything hat occupies space and has mass.
∙ Mass = measure of the quantity of matter that an object has.
∙ Matter can be classified according to physical state (solid, liquid, or gas) or composition. ∙ Pure substances can be elements or compounds.
∙ Elements: consist of only 1 type of atom and cannot be further simplified. (ex; sodium, helium).
∙ Compounds: combination of 2 or more elements, can be further reduced. (ex: pure water). ∙ There are 118 recorded elements on the periodic table.
∙ Elements have a symbol, name, atomic number, and atomic weight. ∙ Mixtures: a combination of 2 or more pure substances.
If you want to learn more check out What is the cognitive model for?
Homogenous mixture: uniform composition throughout. Don't forget about the age old question of What are the different types of earth crust?
Don't forget about the age old question of What is the first known written law codes in history?
Don't forget about the age old question of Spanish built more than how many catholic churches?
Heterogenous mixture: non uniform composition. (can be separated by filtration). When separated the components of both mixtures yields pure substances. ∙ Mixtures can be separated based on physical properties. Such as; density, boiling point, state of matter, vapor pressure, magnetism, solubility, etc.
∙ Chemical compounds are composed of two or more elements
∙ All compounds are made of molecules or ions.
Ionic compounds: composed of ions.
Chemical compounds: compounds have different properties than elements. Fixed composition ratio: definite % composition by mass. If you want to learn more check out Hyper polarization is caused by what?
Example: Iron pyrite FeS2 46.55% iron and 53.45% sulfur.
∙ Physical properties: characteristics that can be evaluated without changing the composition: color, odor, density, volume, hardness, melting point, etc. ∙ Extensive properties: mass; depend upon the amount of the substance. ∙ Intensive properties: density; independent of the sample size. Don't forget about the age old question of Aside from selling and financing, what are the other purposes of marketing?
∙ Density: ratio of ass to volume.
Example) The density of liquid mercury at 20 degrees Celsius is 13.570 g/ml. What is the mass of 107.3 mL of mercury?
13.570g / mL X 107.3mL = 1456.06 mL = 1.456 X 10^3 g.
Example) The density of acetic acid at 20 degrees Celsius is 1.05 g/mL. What volume is needed to have 275 g of acetic acid?
275g/ml (ml/ 1.05g) = 261.90 = 262 mL
∙ Physical change: a change that does not entail any change in the chemical composition. ∙ Chemical change: one or more substances are transformed into one or more NEW substances.
Involves change in chemical composition.
Results in a change in composition or structure.
Chemical reaction: reactants products.
The reactants and products are VERY different.
Example) iron metal reacts with oxygen to form rust (iron oxide). ∙ Tools of quantitative chemistry:
Measurements of length, mass, volume, time, speed, temperature, etc. Each measurement has (1) a number, (2) a label/unit, (3) a name. ∙ Qualitative observations:
no numbers involved, only physical appearance.
Example: color, large/ small, hot/cold, smell, etc.
∙ Units of measurement: Systeme International d’ Unites (SI modified metric system). Mass; unit = kilogram (kg)
Length; unit = meter (m)
Time; unit= second (s)
Temperature; unit = Kelvin (K)
Amount of substance; unit= mole (mol)
** VOLUME IS NOT A BASE UNIT ***
∙ Length: (m)
1000m= 10^3 m = 1 kilometer (km)
10 ^2 m = 1 cm
10 ^3 = 1 mm
0.10 m= decimeter
100 cm = 1 m
∙ Size of molecules are in the nanometer range (nm)
10^9 m= nanometer
Giga (G); 10^9 billion.
Mega (M); 10^6 million
Kilo (k); 10^3 thousand
Deci (d); 10 ^1 tenth
Centi(c); 10^2 one hundredth
Milli (m); 10^3 one thousandth
Micro (/u * fancy u); 10^6 one millionth
Nano (n); 10^9 one billionth
Pico (p); 10^12
Femto (f); 10^15
∙ Volume: NOT a base unit.
Cubic meter is a volume unit m^3.
M = length
M^3 = Volume
Common units: L, mL, cc.
1 cm^3= 1 mL
∙ Mass: base unit = kg.
250 mL of water= 250 g of water= 1 cup.
250kg = 250 L = 1 tub.
∙ Time: base unit= second.
For smaller time frames: milliseconds (ms), microseconds (*fancy u s), nanoseconds (ns), etc.
For more than a second: minutes, hours, days, weeks, etc.
∙ Temperature: kelvin or Celsius.
Kelvin= Celsius + 273.15
Example: The thermostat is set to 78 degrees F, which converts to about 26 degrees C. What is the temperature in Kelvin? 26+273.15= 299.15 K.
∙ Dimensional analysis:
Example) 325 mg aspirin pill. Transform to grams.
325 mg (1g / 1000mg) = 0.325 g of aspirin.
Example) How many inches in ¼ mile?
¼ mile = 0.250 mile
1 mile = 5280 ft.
12 in = 1 ft.
0.250 (5280 ft/ 1 mile) (12 in / 1 ft) = 15,840 inches.
Example) If the speed of a molecule of oxygen as 25 degrees Celsius is 490 m/s and 1 mile is 1.609km, what is the speed of oxygen in mph.
490m/s (1km/ 1000m) (1 mile/ 1.609km) (60 s / 1 min) (60 min/ 1 hr) = 1100 mph ∙ Atomic composition:
Atoms are mostly empty space.
At the center are protons, neutrons, and electrons.
Protons = positive charge
Neutrons = no charge
Electrons = negative charge
The number of protons is always an integer.
∙ Atomic number: (Z) = number of protons.
** in neutral atoms the number of electrons = the number of protons. All atoms of the same element have the same number of protons.
(Z) defines the element.
Example) H=1, Cr= 24, F= 9, Cu= 29
∙ Mass number: (A) = protons + neutrons
Can determine number of neutrons from mass number.
∙ Neutrons = (Z) – (A)
Existence of isotopes means all atoms of an element are NOT exactly the same. All atoms of the same element have the same number of protons.
Atoms of the same element can have different numbers of neutrons in the nucleus. Same protons, different neutrons, different mass number.
Example) C12 = 6 Protons, 6 neutrons, (A)= 12.
C13 = 6 protons, 7 neutrons, (A) = 13.
Example) Hydrogen isotopes: Protium: 1 proton, 0 neutron.
Deuterium: 1 proton, 1 neutron.
Tritium: 1 proton, 2 neutrons.
Cobalt60 used in medical imaging, stable isotope, (Z) 27 protons, 33 neutrons. Iodine isotope: iodine131, (Z)53, 78 neutrons.
Technecium99, medical imaging, (Z) 93, (A)99, 56 neutrons.
Magnesium has 3 stable isotopes: Mg 24, Mg 25, Mg26. All have 12 protons. ∙ Atomic weight: = average of a sample of an element.
Average because most elements are mixtures of isotopes.
The standard is based upon carbon – 12 isotope.
∙ Mass spectrometry: the masses of isotopes and their abundances are determined experimentally by mass spectrometry.
Each isotope is represented by a relative abundance.
∙ Isotope Abundance:
Percent abundance= number of atoms of an isotope/ number of atoms of ALL isotopes X 100
∙ Atomic weight (mass):
Atomic weight = % abundance isotope 1/ 100 (mass of isotope 1) + % abundance isotope 2/ 100 (mass of isotope 2)
∙ Periodic table: Dmitri Mendeleev (18341907) developed the table, arranged the elements according to increasing atomic mass.
Used columns b/c properties are periodic; physical properties, chemical formulas, rxns.
Families /groups are vertical.
Rows are horizontal.
Most elements are metals, then there are nonmetals, metalloids, alkali metals, alkaline earth metals, halogens, noble gases, lanthanide & actinide series.
∙ Compounds: combinations of 2 or more elements in definite ratios by mass. The character of each element is lost when forming a compound.
2 types of compounds:
Molecular compounds: smallest unit of a compound that retains the characteristics of the compound. Nonmetal with metal.
Ionic compound: composed of ions; atoms or groups of atoms with a positive or negative electric charge. Held together by (+) () attraction (ionic bond). Have high melting points.
Most dissolve in water.
Removing electrons from an atom produces a cation with a positive charge. Adding electrons to an atom produces an anion with a negative charge. Metals tend to lose electrons cations.
Nonmetals gain electrons anions.
Monatomic anions are a single element and end in suffix ide.
Group 1A= +1 charge.
Group 2A excluding Be = +2 charge.
Group 3A (only Al) = +3 charge.
Group 5A (N &P ) = 3 charge.
Group 6A= 2 charge.
Group 7A = 1 charge.
Transition metals vary in charge.
Boron and Beryllium do NOT form ionic compounds.
∙ Types of formulas to represent compounds:
Molecular formula: describe the composition, #’s and types of each atom. Ex) C2H6O Ethanol.
Condensed formula: indicate how certain atoms are grouped. Ex) CH3CH2OH Ethanol.
Structural formula: shows the connections and bonds. The lines between atoms represent the chemical bonds that hold atoms together in the molecule. Molecular model: a 3D model.
∙ Polyatomic Ions:
A group of atoms tend to stay together an ion contains atoms covalently bound. Polyatomic anions are groups of atoms with a net charge.
∙ Naming the formula for polyatomic ions:
cation comes first, subscripts must produce neutral unit, subscripts must be the smallest set of whole numbers.
Example) Na Na+ = sodium ion.
Example) Ba Ba 2+ = barium ion.
Example) Al Al 3+ = aluminum ion.
Transition elements vary in charge to identity use roman numerals. Example) Cu Cu + = copper (I) ion.
Example) Cu Cu +2 = copper (II) ion.
Example) Fe +2 = iron (II) ion.
Example) Fe+3 = iron (III) ion.
Naming monatomic anions: add suffix ide.
Example) Fl Fluoride.
Example) O 2 = oxide.
Example) N 3 = nitride.
∙ The mole:
The unit used to count things.
It is a really big number used to count really small things such as atoms, molecules, etc.
1 mole is the amount of substance that contains as many particles as there are in 12.0 g of 12C.
1 mol = 6.02 X 10 ^23 things.
Mole; 6.02 X 10^23
∙ Avogadro's Number:
602,000 x Million x Million x Million
6.0221415 x 10^23
∙ Molecular weight vs Molar mass:
Molecular weight: aka molecular/ atomic mass; mass of 1 molecule in amu (1 atom). Molar mass: mass of 1 mole in grams – element or compound
∙ Mass – mole conversion:
1 mol = formula weight (aka molecular/ atomic weight) in grams. Atoms = atomic weight or mass
Molecules= molecular weight or mass
Ions = formula weight.
Moles to mass conversion:
Moles X grams/ 1 mol = grams
Mass to moles conversion:
Grams X 1 mol/ grams = moles
∙ Percentage composition: calculate the mass percent of each element in a compound. (atoms of an element) (atomic weight) / (formula weight of compound) X 100 % Example) What is the composition of Ca (OH) 2?
Formula Weight = 74.1 g/mol
Ca = 40.078
2 H= 2.016
2 O= 32.000
Add ALL THE ELEMENTS ATOMIC WEIGHT TOGETHER TO GET THE FORMULA WEIGHT.
%Ca = (1) (40.0)/ 74.1 X 100%
= 54.08 54% Ca
∙ Percent composition data can be used to calculate an empirical formula. ∙ Empirical vs Molecular formula:
Empirical: show ONLY THE SIMPLEST ration of atoms in the formula. Almost all ionic compounds are represented by their empirical formulas. Molecular: show the actual composition of a molecule.
Example) Ethylene glycol:
Molecular formula = C2H6O2
Condensed formula = HOCH2CH2OH
Empirical formula = CH3O
Main component in anti freeze.
Example) H2O = Empirical & molecular formula.
Example) H2O2 = molecular, the empirical formula = HO
Example) NO2 = empirical & molecular formula.
Example) C6H1206 = molecular, the empirical formula = CH2O
Deriving an empirical formula:
1) Convert mass percent to mass.
2) Convert mass to moles.
3) Find mole ratio.
4) Convert to whole – number ratio of A to B.
∙ Calculate formula from % composition:
Example) 30.4% N + 69.6% O = 100
30.3 g N (1 mol/ 14g) = 2.17 mol N atom.
69.6 g O (1 mol/ 16 g) = 4.35 mol O atom.
4.35/ 2.17 = 2 mol O / 1 mol N = NO2
Example) A compound is found to be 64.83% carbon, 21.59% oxygen, and 13.59% hydrogen. What is the empirical formula for this compound?
(64.83 gC) ( 1 mol C / 12.011 g C) = 5.397 mol C
(21.59 g O) (1 mol O/ 16.00g O) = 1.349 mol O
(13.59 g H) (1 mol H/ 1.007 g H) = 13.48 mol H
Divide each amount of mol by the lowest mol = 1.349
5.397 mol C/ 1.349 mol O = 4.001 mol C.
13.48 mol H/ 1.349 mol O= 9.992 mol H
1.349 mol O / 1.349 mol O= 1.000
The empirical formula is C4H10O
∙ What if the numbers are not whole numbers?
Multiple each ratio by a whole number to get an even number formula subscripts. ∙ For some compounds, the molecular formula is a multiple of the empirical formula. ∙ Molar formula (g/mol) / empirical formula mass (g/mol) = n
Molar mass / empirical formula mass = n
∙ Determining molecular formula:
The empirical formula weight for C4H10O is 74.12 g/mol
In a separate experiment the molar mass of the compound was determined to be 222.1 g/mol.
What is the molecular formula?
222.1g/mol / 74.12 g/mol = 2.996 = 3
(C4H10O) (3) = C12H30O3
Points to consider:
∙ Nicotine contains 74.0 % C, 8.65% H, 17.35% N. What is the empirical formula? 74. 0 g C ( 1 mol C/ 12.01 g) = 6.16 mol C
8.65 g H ( 1 mol H/ 1.00 g) = 8.58 mol H
17.35 g N (1 mol N/ 14.00g) = 1.24 mol N
Divide each answer by the lowest number of atoms : 1.24 mol N
6.16 mol C/ 1.24 mol N= 4.97 = 5 mol C
8.58 mol H/ 1.25 mol N= 6.92 = 7 mol H
1.24 mol N/ 1.24 mol N= 1.000 mol N
Formula = C5H7N
∙ What elements belong to the alkali metals?
Li, Na, K, Rb, Cs, Fr.
∙ Physical properties of bromine: freezes at 7.2 Celsius, reddish liquid, a sample has a volume of 25.0 mL.
∙ Convert 183 nm to cm.
1.83 X 10*5 cm.
∙ If you entered a 50 km race how far would this be in miles? Use the conversion 5,280 ft = 1 mile, 12 in = 1 ft, 2.54 cm= 1 in.
∙ Calcium 40 and Chlorine 37 are alike because both have?
The same number of neutrons.
∙ An 11.93 g sample of Na contains about ____ atoms?
11.93gNa (1 mol Na/ 23gNa) (6.02 x 10^23 atoms Na/ 1 mol Na) = 7.1819 x 10 ^24/ 23 = 3.123 = 3.1 x 10^23.
∙ With a metal, nitrogen forms ____ with a charge of ____ .
None. Nitrogen does not form ions.
∙ Formula: Al2 (NO3)3
∙ Name & formulas:
KClO4: potassium perchlorate
Mg(OH)2: magnesium hydroxide
CaHPO4: calcium hydrogen phosphate
MgSO3: magnesium sulfite
∙ Name K2S : potassium sulfide
∙ Name Cr(NO3)2: chromium (II) nitrate
∙ Name Cl2O: chlorine oxide
∙ What is the balanced chemical equation for the complete combustion of benzoic acid, C6H5CO2H, to form carbon dioxide water?
2 C6H5CO2H(s) + 15 O2 (g) 14 CO2 (g) + 6H2O (g)
∙ Which of the following compounds are soluble in water: NH4NO3, Fe2S3, CuCO3, and SrCl2?
NH4NO3 and SrCl2
∙ Write the equation for the reaction that occurs when the hydrogen sulfate ion behaves as a Bronsted Lowry acid in water?
Bronsted Lowry acid: “It donates protons in the form of a hydrogen ion (H+). This is reinforced by the definition of an acid, which is a solution that has an excess of hydrogen ions (H+). The BronstedLowry base, on the other hand, is a solution that accepts protons.” (https://study.com/academy/lesson/bronstedlowryaciddefinition examples.html )
HSO4 –(aq) + H2O (l) SO4 2 (aq) + H3O + (aq)
∙ When a solution of solver fluoride is added to a solution of magnesium sulfate, both silver sulfate and magnesium fluoride precipitate.
∙ Alkaline earth metals:
Be, Mg, Ca, Sr, Ba, Ra.
F, Cl, Br, I, At.
∙ Nobel gases:
He, Ne, Ar, Kr, Xe, Rn.