chem exam 2 :3

Aufbau Principle

electrons occupy lowest energy orbital when in ground state

What columns are s-orbitals in?

1-2 (including helium)

2 ways to describe occupied orbitals

Electron configurations and orbital diagrams

Electron configurations

Tell energy level (n), orbital level (s,p,d,f, uses L), and number of electrons in each energy+orbital block (written as superscripts). Does not convey electron spin, all orbitals included

Orbital diagrams

Orbitals drawn as lines/boxes with labels!!! Electrons drawn as arrows (full or half). Indicates electron spin

2 properties followed when electrons fill orbitals

The Pauli Exclusion Principle,

The Pauli Exclusion Principle

Two electrons in an atom can’t have the same 4 quantum numbers. 2 electrons max in 1 orbital, must have opposing spins

Hund’s Rule

When filling degenerate orbitals, electrons fill them one time first, must have parallel spins

• Lose the least energy

Degenerate orbitals

Same energy levels

Aufbau equation relationship w/ hund’s rule

Hund’s rule: application of aufbau’s principle. Electron minimizes repulsions +energy by not sharing an orbital until forced. Parallel spins (in degenerate orbitals) can exchange positions without any electrons lost → stabilizing effect.

Where do gained electrons go?

Go to lowest energy level

Where do you remove an electron?

Highest energy level available

Idk what to name this card

All orbitals have same energy in the same level

Two types of electrons

Valence and core electrons

Valence electrons

Responsible for chemical bonding + atom ionization

Main group elements

S+P orbitals

Where are valence electrons in main group elements?

In highest/outermost n level

Where are valence electrons in transition metals?

In outermost s (E-level n) and d orbitals (E level n-1)

Core electrons

All electrons not valence electrons

When can you use abbreviated electron configurations?

Elements after neon

Coulomb’s Law

gives the P.E. btwn two charged particles based on their charges and distance from eachother

Coulomb’s Law formula

Charges: (q1, qz)

Distance: r

When is potential energy positive in Coulumb’s law?

When charges are similar

• Shows repulsion

• More energy needed to hold position as distance decreases

When is potential energy negative in Coulumb’s law?

When charges are opposite

• Shows attraction

• More attraction as r decreases

How are PE and distance related?

Closer together = higher PE

Farther = less

How is PE affected when value of charges increases

PE increases with magnitude of charges

Shielding

electrons attracted to nucleus, repulsed by other electrons. Affects valence electron attraction, they are shielded by core electrons btwn valence electrons and nucleus.

• Makes valence electrons easier to remove/add

• Valence electrons don’t shield each other.

Effective Nuclear Charge (Zeff)

The attraction to the nucleus a valence electron actually experiences considering shielding done by core electrons

Zeff formula

Periodic trend of Zeff

Increases as you go to the right

First ionization energy periodic trend

decreases as you go to the right

Orbital Penetration

Attraction to nucleus a valence electron would experience if it was closer to nucleus than core electron

• No longer shielded by core electrons

• Zeff: no subtraction of core electrons

Atomic Radius periodic trend

Increases down a column, decreases left to right (because of Zeff, stronger attraction decreases radii)

Atomic Radius

a set of average bonding radii of an element

• Always smaller than nonbonding radius

• Approx. bond length of (z) covalently bonded atoms is the sum of their atomic radii

Nonbonding radius

½ the distance between two nucli in adjacent atoms in a solid state

Bonding radius

Metals: ½ the distance between two nuclei of atoms next to each other in a crystal

Nonmetals: ½ the distance between two nuclei of bonded atoms

Nonbonding radius formula

Atomic radius trends for main group elements

Down a column: radius increases (n increases, larger orbitals occupied)

Across a row: Radius decrease because Zeff

Atomic radius trends for transition metals

Down first two rows of a column: Radius increases

Across a row: Decreases for first few elements, then becomes constant

Ionic Radii periodic trend

Cations: Lose electrons, stronger attraction with other electrons, cations much smaller than neutral atoms. No change in protons

Anions: Gain electrons, Strong repulsion with valence electrons, anions larger than neutral atom. No change in Zeff

Ionization Energy periodic trend

Across a row: IE generally increases (stronger attraction to nucleus)

Down a Column: IE decreases (n increases → valence less attracted)

Ionization Energy

Energy required to remove an electron from atom in a gaseous state

• PE always positive, need energy to overcome attraction w/ proton and electron

• Ie = sum of two IE of elements

• Step-wise format

IE exceptions for going across a row

Pi<s, due to less orbital penetration of p-orbitals into the nuclear region, higher E level for p orbitals (look at spin diagram, paired spins)

p4<p3 because p4 elements have a higher E paired electron removed

Other IE trends

Electron Affinity (EA)

The E change of an atom/ion associated with gaining an electron in the gaseous state (anion stuff)

• Energy released because new attraction, PE lowered

Electron affinity Periodic Trend

Down a column: No definite trend for all families

• Relatively more + for alkali metals

  • Added electrons goes into higher n level orbital that is farther from nucleus

Across a row: Generally becomes more negative (more energy released)

• Added electrons, higher Zeff going across period

• Nonmetals obtain noble gas electron config. when electrons added → most negative EA

Exception for Electron affinity when going across a row

p³>p²: Less to add an electron to elements in column starting with N because p-orbitals are half-filled, causing the added electrons to be spin paired (forced to share orbital)

Metallic characteristics

• Good conductors of heat, electricity

• Malleable, ductile

• Tend to lose electrons (Less IE) (cations)

Nonmetals

• Poor conductors

• Tend to gain electrons (less EA, more negative) (anions)

Metallic character periodic trend

Down a column - metal character increases

Across a row - metal character decreases

Magnitude

The size of a value

Does magnitude of PE increase or decrease as values of charges increase?

Increase

Are sublevels of n of s,p,d orbitals degenerate?

No

Cr column electron configuration

• CR and Mo prefer half-filled d-orbital, leaving a half-filled s-orbitals

  • D and p half-filled orbitals typically stable

  • Because change in energy between 4d orbitals is less than the repulsion of energy paired in s-orbital

Cu column electron configuration

(Cu, Ag, Au) Prefer full d-orbital and half-filled s orbital

• Difference in energy btwn s and d orbitals are still smaller than repulsions of electrons paired in d orbital vs. s orbital

Transition metal electron configurations

Fill n s-orbitals first, then n-1 d orbitals second

• Remember hund’s rule

Magnetic properties of atoms + ions

• Magnetic field created with moving electric charges → result of electron spin unless it orbital is spin-paired (full)

2 types of magnetic atoms/ions

Paramagnetic, Diamagnetic

Paramagnetic

Atom or ion w/ any unpaired electron

• Result: particle is attracted to an external magnetic field

• Electrons not cancelled out by another pair

Diamagnetic

An atom/ion as all electrons paired

• Particle is not attracted to external magnetic field

• Electrons cancelled out by another pair

Copper electron configuration

4s^1 3d^10

Which electron do you remove in a transition metal?

The electron with the highest n value (disregard energy values)

Molecules

Any group of 2 or more atoms bonded together

Compounds

Type of molecule, made of 2 or more elements

Homonuclear

Same atom

Column 1 category of periodic table

Alkali metals (-1)

Column 2 category of periodic table

Alkaline Earth metals

Column 18 category of periodic table

Noble Gases (-1)

Column 17 category of periodic table

Halogens

Column 16 category of periodic table

Chalcogens (-2)

Chemical bonds

a force that holds atoms together

• Occurs due to attraction w/ charged particles

  • Valence electron of 1 atom attracted to nucleus of 1 atom

When does chemical bonding occur?

Only forms if attractive forces > repulsive forces

Nonmetal characteristics

Varied properties

• Conduction of heat/electricity: Poor

• Tend to gain electrons

Metalloid characteristics

Distinct properties

• Conduction of heat/electricity: Semi

Metal characteristics

• Conduction of heat/electricity: Good

• Tend to lose electrons

2 types of Chemical bonds

Ionic, Covalent

Ionic bonds

Metal+nonmetal bond

• Transfer of electrons

• Opposite charged electrons attract each other (columb’s law)

Covalent bonds

2 or more nonmetals

• Share electrons

• Result of difficulty of removing electrons from nonmetals (high IE)

• Attracted to both atom’s nuclei (lower PE, columbs law)

Where are nonmetals and metalloids on periodic table?

Are metals cations or anions?

Cations

Are nonmetals cations or anions?

Anions

Ionic compound

Composed of a lattice (3D array, several atoms) of alternating cations and anions in a solid state

Formula unit

• Smallest electrically neutral collection of ions

• Ex. CaCl2 (Ca+2 and Cl-, have two Cl- to make the compound neutral)

Molecular Compounds

Discrete, neutral molecules

• No ions !!!!!!!!!!!!!!!!!92731456738294576483294857648392

• Represented with chemical formulas

  • Indicates no. of each type of atom for specific compound

Empirical Formula

Give the relative # of atoms of each element (simplest ratio)

Molecular formula

Gives actual # of atoms of each element

Structural Formula

Represents how the atoms are bonded and/or geometry of molecule

Three types of structural formulas

Line formula, condensed structural formula, organic condensed line formulas

Line formula

Represent covalent bonds as lines btwn 2 atoms

• 1 line = 2 electrons

Condensed Structural Formulas

Give chemical formula w/o bond lines

• indicate how atoms are connected w/in the formula

Organic Condensed Line formulas

Depict C atoms as terminal pt./vertice.

• Omit H bonded to C

• C has 4 bonds in total (can be single, double bonds….)

2 types of masses from a molecule’s chemical formula/formula unit

Formula mass, Molar mass

Formula mass (Amu)

avg. mass of molecule

• Add up atomic masses of all elements in compound

Molar Mass (MM)

The mass (g) of 1 mol of a compound as defined by formula unit

• Add up molar mass of all elements in compound (g/mol)

Mass percent composition

(element g/mol you want to find / total MM) x100

Molar relationship w/ elements and compounds

No. of moles of each element in 1 mol of compound

2 types of compounds

Ionic+Molecular

Type I metals

Always know charge (Alkali metals+earth metals, +1 and +2) and Al, Zn, Ag

Zn charge

+2

Ag charge

+1