Chem exam 1

density

mass of certain volume of substance

D = M/V

precision vs. accuracy

Precision: how close together different measurements are

Accuracy: how close measurement is to actual result

celsius to kelvin

T(kelvin) = T(celsius) + 273.15

leading zeros

in front

0.015 (2 sig figs)

captive zeros

in between other numbers

chem 101 (3 sig figs)

trailing zeros

in the back

150 vs 150. vs 150.0 (2sf, 3sf, 4sf)

exact numbers

not measurements

don’t count toward sig figs

scientific notation

11000 = 1.1 × 10^4

same number of significant figures on both sides of each example

sig fig rules: multiplication & division

take lowest number of sig figs

4.56 × 1.4 = 6.38 → 6.4

sig fig rules: addition & subtraction

same number decimal places as least precise

1.3 + 1.225 + 10.45 = 12.975 → 13.0

standardized units

mass: kilogram (kg)

length: meter (m)

time: second (s)

temperature: Kelvin (K)

amt of substance: mole (mol)

electrical charge: Coulomb (C)

metric prefixes (need to know centi, milli, kilo)

• centi = 0.01

• milli = 0.001

• kilo = 1000

mega, kilo, hecto, deka, UNIT, deci, centi, milli, micor, nano

(Mighty king hector died unexpectedly drinking chocolate milk Monday night)

heterogeneous mixture

distinguishable parts

heterogeneous mixture has visibly different parts or regions that are not evenly distributed

homogeneous mixture

indistinguishable parts (solution)

has a uniform composition and appearance throughout

solid

definite shape + volume

liquid

definite volume + variable shape (takes shape of its container)

gas

variable shape and volume (takes both of its container)

pure substance

form of matter that has a constant, uniform composition and distinct, unchanging properties throughout

cannot be separated by any physical means

• separation of saltwater: distillation

atomic theory

Law of the Conservation of Matter

• Mass is neither created or destroyed in chemical process

  • just moved around

  • limits possibilities of chemical reactions

Law of Multiple Proportions

• When two elements form a series of compounds, the ratio of the masses of each elements can always be reduced to a whole number

Nature only allows certain combinations of masses

Oxygen and hydrogen mostly react with the ratio: 8g O / 1g H

proton

positively charged subatomic particle found in nucleus

electron

negatively charged subatomic particlen

neutron

uncharged subatomic particle w/ mass approx. same as protons

isotopes

atoms of the same element that differ in mass

• different number of neutrons

atomic number (Z)

number of protons in nucleus of an atom

mass number (A)

total number of protons and neutrons in an atom

atomic symbols

ion

electrically charged atom, numbers of subatomic particles are not equal

cation vs. anion

ION - atom that has unequal number of protons and electrons

CATION - positively charged ion that forms when an atom/molecule loses one or more electrons

• more protons than electrons

ANION - negatively charged ion that forms when atom/molecule gains one or more electrons

• more electrons than protons

atomic mass

Average corrected for relative abundances of each isotope

• To find the average atomic mass of an element, multiply the mass of each isotope by its decimal abundance and add the results together

mass spectrometry

Mass spectrometers = instruments to determine the mass of substances.

• tells difference b/w isotopes

– Convert molecules into ions.

– Separate ions based on mass/charge ratio.

How to solve mass spectrometry:

A sample of oxygen (O 2 ) gas is isotopically enriched such that 50% of the atoms are 16 O and 50% of the atoms are 18 O. What is the correct mass spectrum for this gas?

1. recognize that O2 is diatomic

2. list the possible molecule masses

   1. 16 + 16 = 32

   2. 16 +18 = 34

   3. 18 + 18 = 36

3. calculate the probability of each combination

   1. 0.5×0.5=0.25

   2. 0.5×0.5=0.25

   3. 2×0.5×0.5=0.50

4. translate to a spectrum

   1. 1:2:1 ratio

ionic compounds

• Metal with nonmetal

• Electron transferred from one atom to another and electrostatic attraction holds compound together

covalent compounds

• Nonmetal with nonmetal

• Electrons spend time between atoms and proton-electron attraction holds atoms together

organic compounds

Carbon containing compounds (excluding carbon oxides)

acids/bases

Species that donate or accept protons (one definition)

naming binary ionic compounds

1. Cation named first and anion second

2. Cation takes name from name of parent element

3. Anion named by taking the root of the element name and adding –ide

4. No Prefixes are used

e.g. KCl = potassium chloride

naming ionic compounds w/ more than one oxidation state

• Transition metal cations usually require a Roman numeral

• Charge on the metal ion must be specified

• Roman numeral indicates the charge of the metal cation

e.g. FeCl₂ = iron(II) chloride

nomenclature prefixes

1 mono-

2 di-

3 tri-

4 tetra-

5 penta-

6 hexa-

7 hepta-

8 octa-

9 nona-

10 deca-

naming covalent compounds

Formed between two nonmetals.

1. First element in the formula (more metallic) is named first, using the full element name.

2. Second element named as if it were an anion (ends in –ide)

3. Prefixes are used to denote the numbers of atoms present.

4. The prefix mono- is never used for naming the first element.

e.g. CO₂ = carbon dioxide

naming binary acids (hydrogen and metallic element, NO oxygen)

If the anion does not contain oxygen, the acid is named with the prefix hydro– and the suffix –ic.

e.g. HCl = hydrochloric acid

naming oxyacids (contain H, O, and at least 1 other element)

If the anion does contain oxygen:

• The suffix –ic is added to the root name if the anion name ends in –ate.

• The suffix –ous is added to the root name if anion name ends in -ite

e.g. HC₂H₃O₂ = acetic acid; H₂SO₃ = sulfurous acid

the mole

1 mole of anything = 6.022 x 1023 units of that thing (Avogadro’s number).

mass percent of an element

mass % = (mass of element in compound / mass of compound) * 100%

chemical formula

Example: benzene

• Empirical formula = CH

  • Simplest whole-number ratio

• Molecular formula = (empirical formula)ₙ [n = integer]

• Molecular formula = C₆H₆ = (CH)₆

  • Actual formula of the compound

determining empirical formula

General rules for Empirical formula from starting from mass percent values:

• Assume 100 g of total compound

• Calc. # of moles of each element

• Divide each value of moles by smallest value

• If all values are whole numbers, the number for each element is the subscript

• If not, multiply by an integer to get whole numbers

determining molecular formula

• Obtain empirical formula

• Calculate mass of empirical formula

• Divide molar mass by mass of empirical formula

• Multiply value by empirical formula

combustion reaction

rapid addition of oxygen to produce carbon dioxide and water

combustion analysis

• Determine the % C and H in a sample

• Assume all C and H transferred into carbon dioxide and water

• Amount of oxygen unknown (often deduced by subtraction of % H and C from 100)

• Often used to find empirical formula

combustion reaction is rapid addition of oxygen to produce carbon dioxide and water

balancing equations

• Number of atoms of each element must be same on both sides of balanced equation

• Subscripts must not be changed to balance equation

• Coefficients can be fractions

tips to balance equations

• if an element is present in just one compound on each

side, balance it first

• balance anything that exists as a free element last

• balance polyatomic ions as a unit

• check when done – same number of atoms, and same total charge (if any) on both sides

Calculating Masses of Reactants and Products in Reactions

1. Balance equation for reaction

2. Convert mass of reactant or product to moles

3. Use balanced equation to set up appropriate mole ratios

4. Use mole ratios to calculate number of moles of desired reactant or product

5. Convert from moles back to grams if required by the problem

In general: grams → moles → moles → gram

limiting reagent

the starting substance that is completely used up first in a chemical reaction, which stops the reaction from continuing and limits how much product can be made

how to determine limiting reagent

1. write balanced equation

2. convert given amounts to moles

3. calculate product yield (smaller amount = LR)

or, balance equation, then take each reactant and find the lowest quantity per coefficient ratio from atoms/molecules/moles

how to find reactant in excess

• Balance the equation: Write the correct balanced chemical equation for the reaction.

• Identify the limiting reactant: Convert all given reactant quantities into moles, then determine which reactant runs out first.

• Calculate the amount used: Use stoichiometry and the mole ratio from the balanced equation to find how much of the excess reactant reacts with the limiting reactant.

• Subtract the used amount: Take your initial starting amount of the excess reactant and subtract the amount consumed.

• Convert units if needed: Change the final leftover value back into grams or your desired unit

percent yield

actual yield = what you get

theoretical yield = what you would get if chemical reaction was strictly followed

(actual yield / theoretical yield) * 100% = percent yield

substance

kind of matter with uniform properties

atom

fundamental unit of a substance

molecule

two or more atoms held in specific shape by attractive forces

elementary substance

substance w/ only atoms of the same atomic number

compound substance

substance with more than one type of atom

chemical formula

elements listed with number of atoms as subscript

diatomic elements

hydrogen (H₂)

nitrogen (N₂)

fluorine (F₂)

oxygen (O₂)

iodine (I₂)

chlorine (Cl₂)

bromine (Br₂)

have no fear of ice cold beer

alkali metals

six chemical elements that make up Group 1 of the periodic table:

• lithium (Li),

• sodium (Na)

• potassium (K)

• rubidium (Rb)

• cesium (Cs)

• francium (Fr)

alkaline earth metals

Six chemical elements that make up Group 2 of the periodic table.

List of Elements

• Beryllium (Be)

• Magnesium (Mg)

• Calcium (Ca)

• Strontium (Sr)

• Barium (Ba)

• Radium (Ra)

halogens

the six non-metallic elements that make up Group 17 of the periodic table.

The Halogen Elements

• Fluorine (F):

• Chlorine (Cl):

• Bromine (Br):

• Iodine (I):

• Astatine (At):

• Tennessine (Ts):

noble gases

seven noble gases are

• helium (He),

• neon (Ne)

• argon (Ar)

• krypton (Kr)

• xenon (Xe)

• radon (Rn)

• oganesson (Og).

They make up Group 18 of the periodic table.

physical properties

observed w/o change in composition of substance

• ex. color, density, melting/boiling points (these r intensive…)

intensive: do not change based on how much of the substance you have

extensive: change depending on amt. of matter present (mass, volume, length)

chemical properties

only evident if substance goes through chemical change

• ex. enthalpy of combustion, flammability, toxicity, stability, oxidation state

chemical change requires new substance before and after

compounds with common names

• water H₂O

• ammonia NH₃

• ozone O₃

writing ionic formulas

• metals tend to lose one or more electrons to form positive ions

• nonmetals tend to gain one or more electrons to form negative ions

• symbol for the metal given first in formula

ex. aluminum and oxygen

Al³⁺ and O²⁻

→ Al₂O₃

stoichiometry

use of coefficients in balanced equations to decide the amount of each reactant that is used, and the amount of each product that is formed