Bohr Model

Bohr Model orbits

-Electrons move in fix orbits around the nucleus

*Based on the hydrogen atom

Bohr Model energy

• Fixed energy

• Electrons stay in those orbits unless they gain energy

• Electrons in orbits closer to nucleus contain less energy than electrons in higher orbits

• Each orbit has certain amount of energy

When can we see light?

• When electrons move from a higher to lower principal energy level

• When they release energy ( loses) therefore it produces light

-Colors from excited gases arise because electrons move between energy states in the atom

Max Planck

• Energy can only be absorbed or released from atoms in certain amounts called quanta

-Light travels in energy packets called photons

-Photon is a quantity of energy emitted

Equation for the energy of one photon

E=H*V

H is constant=6.63 x 10 ^ - 34

V is frequency

What relationship does energy and frequency have?

-It is directly proportional

• As one increases so does the other

Monochromatic

-Radiation composed of one wavelength

Continuous

-Radiation that spans a whole array of different wavelengths

*White light is continuous

Spectroscopy

the process of producing and analyzing spectra

-Important tool for identifying substances

-provides both qualitative and quantative information

Ground State

-The level where the electron is found

• lowest possible energy level for an electron

Excited State

• An energy level higher than the ground state

Atomic emission spectrum

• the patten of frequencies obtained by passing light emitted by atoms of an element in the gaseous state

what is the equation of the speed of electromagnetic wave?

C=V*λ

C is constant = 3.0 x 10^8

λ is wavelength and lambda's number

What is the relationship between wavelength and frequency?

As one increases the other decreases

• Inversely related

Electrons changing orbits

• The closer to the nucleus the bigger the energy change

• When an electron absorbs a specific quantity of energy it "jumps" to a higher level

• When an electron in a higher orbits releases a specific quantity of energy it "jumps" to a lower orbit

Order of wavelengths

radiowave, microwav,e infrared ,visible ( red is the longest wavelength ( ranges from 700 nm-400nm), ultraviolet, x rays, gamma rays

Equation to calculate the change in energy when an electron makes a transition between higher levels

• E=Ef-EI

• Both Ef and Ei contain -2.18 x 10^ -18

If absorbed E is + (positive)

If released/emitted E is - ( negative)

Ionized

-Completely remove from an atom

n= infinity

Quantum Mechanical model

-"Wave" model

• Electrons have quantized energy (electrons are located in specific energy levels)

• In the shell there is still 90% probability of finding an electron

Principal quantum number

-Gives approximate amount of energy of the electron and distance from nucleus

• As increase distance away there is more energy

• Different shapes have different energy levels bece as they further and further away from the nucleus and the electrons are somewhere in that region

Subshell levels

• Specifies the shape of the orbital

• Excited states are unstable (electrons in a higher energy orbital will fall/relax back to lower energy

*The higher the energy level the more sub-levels there is

S sublevel: spherical shaped orbital ( 1 orientation and 2 electrons in total)

P sublevel: dumbell shaped orbital ( 3 different orientaion and 6 electrons in total)

D sublevel: complex shape ( 5 different orientation and 10 electrons in total)

F sublevel: complex shape (with 7 different orientation and 14 electrons in total)

• In order for 2 electrons to occupy the same orbital they must have opposite spins

Different principal energy levels + orbitals

3rd principal energy has: 18 electrons in total

3p, 3s, 3d

4th principal energy has: 32 electrons in total

4s,4p,4d,4f

*f block starts at the 6th level of the other levels

Pauli exclusion principle

• orbitals may hold no more than 2 electrons and they must have opposite spins

• The subshells within a principal shell do not have the same energy ( 4s is lower energy than 3d)

Hund's rule

• When filling orbitals of equal energy ( each subshell), electrons fill them singly first, before pouring with other

• Fill each sublevel first

ex if it is Al ( or has an uneven amount of electorns) fill 2s before 2p

Dual nature of light

Light exhibits both wave properties and particle properties

Debroglie

Matter exhibits both particle properties and wave properties ( electrons have wave properties)

Schrodioeher

• Developed the quantum mechanical mode that uses probability to predict where electrons are located

Heisenberg's uncertainty principal

• It is impossible to know both velocity ( speed+direction) and the location an electron at the same time ( you can only know one)

Electron cloud

• Where electrons can be found

• represents position where there is probability in finding an electron

Orbital vs orbit

-Orbital is the region in space where there is a high probability of finding an electron based on mathamatical calculations

-Orbits: Bohr though electrons move in a smooth path around the nucleus

Electron configuration

the arrangement of electrons around the nucleus, providing insight on element's properties

Aufbau principle

-Electrons occupy orbitals of lowest energy first

Valence electrons

All of the electrons in the highest energy level in an atom

• Count all of the electrons in the highest energy level

Isoelectron

Having identical electron configuration ( ex: Ne and O-2)

Ocetet Rule

• Atoms will try to atttain a more stable configuration by gaining or losing electrons in order to be isoelectron ( the same as) a noble gas

Transition metals and losing electrons

• Form different charges when forming ions

• They'll always lose the highest s electrons first then d electrons

Stability of atom

• Full valence level ( noble gases) are the most stable

• half filled valance subleve is the second most stable

• The least stable is any other configuration

Electron density

• The electron density directly depends on the shape of orbital

p block metals and losing electrons

When in the p block, you lose the p electrons first

excited state configuration

Electrons will jump to the highest energy level

• When energy is released they jump to the lower sublevel

• In electron configuration the way to tell is if there’s an electron that doesn’t belong in the sublevel ( skips one)

Exceptions configurations

Ag: [Kr] 5s1 4d10

Cu:[Ar] 4s1 3d10

Cr:[Ar] 4s1 3d5

Mo:[Kr]5s1 4d5

The periodic law

when arranged by increasing atomic number the chemical elements display regular and repeating patterns of chemical and physical properties ( density, reactivity, states of matter)

periodic trends

how a specific properties change as you move along the table

Shielding effect

-inner electrons shield the outer valence electrons

*shielding effect remains the same across the period but increases down a group because you are adding another sublevel

• This cause the atomic radius to increase as you go down

• this decreases the ionization energy as you go down a group

Ions and how they affect size

• Cations are always smaller than the neutral atom ( protons have more force, more attraction, smaller distance)

• Anions are always larger than the neutral atom ( extra electron is shielding making the distance longer

• The size changes are due to either greater repulsion of more electrons or increased proton attraction of fewer remaining electrons

Ionization energy

• The amount of energy required to remove a valence electron from an atom

• The weaker the attraction the less energy required to remove it ( decreases as you go down a group)

• The greater the attraction the more energy required ( increases as ou go across a period)

Electronegativity

• The ability of an atom to attract (its own) electrons that are shared with another atom in a covalent chemical bond

• the further down a group the less electronegativity there is ( increases as you go across because there is a greater attraction between protons and electrons)

• Scale from 0 to 4.0

• bonds form because atoms want to be stable ( noble gases don't form bonds because they are the most stable)

Metallic character

Metals who lose their electrons the easiest

• biggest atoms have weakest attractions to electron

• increases as you go down a group and decreases as you go across the period

who invented the periodic table

Dmitri Mendeleev made the periodic table increasing by atomic mass.

Characteristics of Ionic bonds vs Covalent bonds

Ionic: Strong, are solids,have higher melting points, good electric conductors/ electrolytes, and soluable in water.

Covalent: can be solids, liquids, gas, melting point and boiling point is lower, and are sometimes soluable in water.