The Mole and Molar Masses

• The mole

  • The mol is a counting unit in chemistry

  • 6.022×10²³ units

  • also called alvagaldos number

  • defined as the number equal to the number of carbon atoms in exactly 12 grams of pure carbon -12

    • meaning that the whole period table is based on carbon -12

• Using the mole In calculations

  • atomic value (amu) found on the table tells us how many grams are in one mole of that element

  • Ex:

    • how many moles of copper are. in a penny with a mass of 3.14g

      • start with 3.14 g- turn into moles with the molar mass

      • to turn into atoms → turn moles into the counting unit

• Molar Mass

  • mass in grams of one mol of the compound (g/mol)

  • how to find

    • 1. account for every mole of each compound found in the element

    • turn the moles into grams using the average atomic masses given on the table

Mass Spectroscopy and Average Atomic Mass

• atomic masses are based on carbon-12 as the standard

  • c-12 will have exactly 12 amu

• Mass Specroscopy

  • an instrumental method that identifies the chemical constitution of a substance by separating it into gasious ions and based on mass and chchrge

    • used to find the relative abundance and the atomic/molar mass of an unknown sample

  • y-axis is the relative abundance of particles

  • x-axis is the m/z or mass over charge

    • usually one electron is taken away so the x axis just represents the mass

  • Steps:

    • vaporization - substances must be gasses

    • ionization - knocking off electrons (usually just one)

    • Accelleration - to have equal KE

    • Deflection - a magnetic field will deflect them, the lighter they are, the more they are deflected, leading them to a sharper turn

    • Detection - detector plate detects ions

• Average Atomic Mass - the weighted average of the masses of the naturally occurring isotopes for that element

  • are located on the periodic table

  • how to calculate average atomic mass when given percents and mass

    • multiply the percents in decimals times the atomic mass

  • How to use spectrometer graph

    • it will provide relative abundances and mass number, calculate based on that

Percent Composition of Compounds and Determining the Formula of a Compound

• there are 2 ways to describe the composition of a compounds

  • Percent Composition by mass

    • can be determined by comparing the mass of each element present in 1 mole of the compound to the total mass of 1 mole of the compound

    • mass percent of the element = (mass of element in 1 mole of compound)/ (mass of 1 mole of compound) times 100%

    • example:

      Find the percent composition from formula of N2O5

      1. we know that one mole of dinitrogen pentoxide has 2 moles of nitrogen and 5 moles of oxygen

      2. we find the mass of each element using stoichiometry to turn moles of the element to grams

      3. then we find the total mass of the compound

      4. then we use the formula for the mass percent of the element found above

      5. NO SIG FIGS because moles are exact.

      Example: what if we were given a substance in grams that yielded elements in grams

      1. just divide the elements in grams over the substance before in grams

    • Always make sure everything is in grams to find percent composition in grams

    • Another Situation:

  • Formula of a Compound

    • Molecular Formula: is the exact formula that gives the type of atoms and the number, this is only for nonmetals

    • Formula unit: for ionic compounds, is always empirical, so it. is the lowest ratio

    • Emperical formula: simplest ratio of a formula and can be calculated with the percent composition’s

    • how to get empirical formula from percent composition

      ex: Ba: 69.58%, C: 6.090%, O: 24.32%

      1. Turn the percents into grams but pretend it was from 100 grams

      2. then with the grams of each element, turn into moles through stociometry

      3. then divide by the smallest moles

      4. multiply all until you get 0.9 or 0.1 from a full mol

      5. that is how many moles of each element are in the empirical formula

      6. what if there. is a missing percent? then you just subtract. the ones given from 100%too find the missing one

      7. what if I was given grams from a sample rather than percents? then turn grams into moles. and continue on with the steps

    • Always turn to moles

    • How to turn empirical into the molecular formula

      • find the molar mass of the empirical formula

      • divide the molar mass (given) over the empirical formula mass

      • the value given will be the number to multiply by all the empirical formula to get the correct molecular formula

  • Emperical Formula from Combustion

    • combustion is the burning of a hydrocarbon in presence of oxygen and → CO2 and H2O

    • ex: given grams of original sample, given yielded CO2 and H2O in g

      1. turn grams of CO2 and H2O into moles of C and H

         1. g → moles CO2 → 1molC/1mol CO2 → moles C

      2. Then find grams of C and H and subtract that from the sample to find mass of O

      3. mass of O → moles

      4. Dicide by smallest mole, then multiply to get full mole number

      5. if given the actual molar mass → Molar mass/emperical mass → coefficient

Polyelectronic Atoms, The Aufbau Principle, and the Periodic Table

• Polyelectronic Atoms - is an atom having more than one electron

• Aufbau Principles - an electron occupies the lowest energy orbital that can receive it

  • 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s

  • use the table

• Perioid Table Blocks

  • Tells us what the last electron filled is

• Electron Confifuration - is how the elcetrons are distributed among the various atomic orbitals in an atom

  • 1s² (the 1 is the pinricpal quantum number), S is the type of subshell? and the exponent is the number of e-

  • Shorthand Notation - find the closest noble gas that has less electrons

    • mark the noble gas with []

    • resume the electron config

• Orbital Diagram

  

• Hunds Rule: when filling orbitals with electrons, each orbital in a subshell should be singly occupied before any orbital is doubly occupied

• Pauli's Exclusion Principle states that no two electrons in the same atom can have identical values for all four of their quantum numbers.

  • meaning that no more than two electrons can occupy the same orbital and (2) two electrons in the same orbital must have opposite spins

• Exceptions to Audbau Principle:

  • d and f orbitals require a lot fo energy

  • d^4 and d^9 exceptions

    • they are one electron short of being held full

    • to become stable (less energy), they take an e- from the closest s orbital

• Ions in Electron Config

  • Electrons come out from the highest energy level not the orbital

  • ex:

    • Atom: [Kr] 5s2 4d10 5p2

      Sn+2 ion: [Kr] 5s2 4d10

      Sn+4 ion: [Kr] 4d10

Photoelectron Spectroscopy

• used to understand atomic structure, electron config, ionization energy, and periodic trends

• What is PES?

  • the photoelectron spectrophometers use high energy radiation photon to eject (ionize?) electrons from an atom

    • bc electrons are at different energy levels, they require different amounts of energy

      • valence easy, inner core are harder

• What is measured?

  • the photon energy removes the electron, the left over KE energy determines how fast the electron is moving

    • ifthe electron is slow, it means it required a lot of ionization energy to remove the electron

    • if. theelctron is fast, a lot of KE, it means little energy was required to remove the electron

• Reading the Graphs

  • Y axis - signal intensity, number of electrons or relative number of electrons

  • X- axis: represents the ionization energy or the binding energy

    • it is. backward x axis

    • to the left, means lots of energy to remove e-, probably inner elections

    • to the right, it is little energy to remove the valence electrons

Periodic Trends

• Periodic Table

  • was originally constructed to represent the patterns observed in the properties of elements

  • the first chemist was Dobereiner with a triad model

  • John Newlands Octave model

  • Meyer and Mendeleev - atomic mass

    • Mendeleev was given the most credit because he was able. to predict the existence and properties of unknown elements like atomic masses

  • current periodic table is made by Henry mosley, based on the atomic number

• Valence Electrons - are the outermost energy. level electrons

  • elements with similar valence configuration show similar chemical behavior

  • in the main group representative elements, the groups have the same electron config

  • predicting the valence electron config of transition metals

• Core Electrons are the inner electrons

• Columbic Attraction: the positive and negative attraction

• Effective Nuclear Charge is the pull that an electron “feels” from the nuclear

  • (Z_eff) = # protons - # core electrons

  • the closer an electron is to the nuclear, the more pull it feels

  • as the effective nuclear charge increases, the electron cloud is tighter

• Ionization Energy - the energy required to remove an electron from a gasous atom or ion

  • X (g) + energy → X⁺(g) +e^-

  • the highest energy electron (the one bound least tight) is removed first

  • when removing electrons from ions, it becomes harder because there is a larger proton to electron ratio

  • when removing from a full shell, the energy is greater by a lot

  • left to right first ionization increases because more protons in the nuclear (higher effective nuclear charge) because more protonts and smaller radius

  • first ionization energy decreases in going down a group because the electrons removed are farther because in a different energy level from the nucleus, lower columbic attraction “feel lesss of nuclear force”

• Atomic Radius

  • the radius is defined as hald the distance between the nucli in a molecule consisting of identical atoms

  • decrease when going left to right: because the increasing effective nuclear charge and the valance electrons are drawn closer

  • increases down a group: because of the increases in the obrital size bc more level

• Size of Ions

  • Negative ions are always larger than the atoms of what they are formed

    • this is because of electron repulsions between them increase and electrons push apart and occupy more volume

    • and lower protein pull compared

  • Positive Ions are always smaller than the atoms from they are formed

    • because the electrons are removed from valence electron and electron repulsions decrease and the elcetrons pulled together close

    • also because the electrons pull them closer

• Electronegativity - is a measure of the ability of an atom in a chemical compound to attract electrons

  • valence electrons hold atoms together in chemical compounds

  • some compounds have valance electrons concentrated closer to one atom than another and this uneen concentration affects the properties of a compound

  • increases across each period

  • decreases or stays the same down the period

  • most electronegativ is flourine

• Electron Affinity - the amount of energy involved when an electron is accepted by a gasous atom. tofodm a negative ion

  • “neutral atoms likelihood of gaining an electron”

  • Metals don’t want to gain electron to be stable so it might be endothermic or very low

  • Nonmetals want an electron

  • Halogens require energy to receive

  • increases left. toright

  • decreases as you go down