Ch. 4 Arrangement of Electrons in Atoms 4.1 The Development of a New Atomic Model

Light Before 1900, scientists thought that light behaved only as wave discovered that also has particle-like characteristics

Light as a Wave electromagnetic radiation: form of energy that acts as a wave as it travels includes: visible light, X rays, ultraviolet and infrared light, microwaves, and radio waves All forms are combined to form electromagnetic spectrum

Light as a Wave

Light as a Wave all form of EM radiation travel at a speed of 3.0 x 108 m/s in a vacuum it has a repetitive motion wavelength: (λ) distance between points on adjacent waves; in nm (109nm = 1m) frequency: (ν) number of waves that passes a point in a second, in waves/second Inversely proportional!

Photoelectric Effect when light is shone on a piece of metal, electrons can be emitted no electrons were emitted if the light’s frequency was below a certain value scientists could not explain this with their classical theories of light Ex: coin-operated sift drink machine

Photoelectric Effect Max Planck: a German physicist suggested that an object emits energy in the form of small packets of energy called quanta quantum- the minimum amount of energy that can be gained or lost by an atom Planck’s constant (h): 6.626 x 10-34 J*s

Photoelectric Effect Einstein added on to Planck’s theory in 1905 suggested that light can be viewed as stream of particles photon- particle of EM radiation having no mass and carrying one quantum of energy energy of photon depends on frequency

Photoelectric Effect EM radiation can only be absorbed by matter in whole numbers of photons when metal is hit by light, an electron must absorb a certain minimum amount of energy to knock the electron loose this minimum energy is created by a minimum frequency since electrons in different metal atoms are bound more or less tightly, then they require more or less energy

H Line-Emission Spectrum ground state- lowest energy state of an atom excited state- when an atom has higher potential energy than it has at ground state line-emission spectrum- series of wavelengths of light created when visible portion of light from excited atoms is shined through a prism

H Line-Emission Spectrum scientists using classical theory expected atoms to be excited by whatever energy they absorbed continuous spectrum- emission of continuous range of frequencies of EM radiation

H Line-Emission Spectrum Why had hydrogen atoms only given off specific frequencies of light? current Quantum Theory attempts to explain this using a new theory of atom

H Line-Emission Spectrum when an excited atom falls back to ground state, it emits photon of radiation the photon is equal to the difference in energy of the original and final states of atom since only certain frequencies are emitted, the differences between the states must be constant

Bohr Model created by Niels Bohr (Danish physicist) in 1913 linked atom’s electron with emission spectrum electron can circle nucleus in certain paths, in which it has a certain amount of energy

Bohr Model Can gain energy by moving to a higher rung on ladder Can lose energy by moving to lower rung on ladder Cannot gain or lose while on same rung of ladder

Bohr Model a photon is released that has an energy equal to the difference between the initial and final energy orbits

Bohr Model problems: did not work for other atoms did not explain chemical behavior of atoms

Introduction to Quantum Theory Quantum Theory- describes mathematically the wave properties of electrons

Electrons as Waves In 1924, Louis de Broglie (French scientist) suggested the way quantized electrons orbit the nucleus is similar to behavior of wave electrons can be seen as waves confined to the space around a nucleus waves could only be certain frequencies since electrons can only have certain amounts of energy

Electrons as Waves shows that anything with both mass and velocity has a corresponding wavelength

Uncertainty Principle In 1927 by Werner Heisenberg (German theoretical physicist) electrons can only be detected by their interaction with photons any attempt to locate a specific electron with a photon knocks the electron off course Heisenberg Uncertainty Principle- it is impossible to know both the position and velocity of an electron

Schrödinger Wave Equation In 1926, Erwin Schrödinger (Austrian physicist) his equation proved that electron energies are quantized only waves of specific energies provided solutions to his equation solutions to his equation are called wave functions

Schrödinger Wave Equation wave functions give only the probability of finding an electron in a certain location orbital- 3D area around a nucleus that has a high probability of containing an electron orbitals have different shapes and sizes

Quantum Numbers specify the properties of atomic orbitals and of electrons in orbitals the first three numbers come from the Schrödinger equation and describe: main energy level shape orientation 4th describes state of electron

1st Quantum Number Principal Quantum Number: n main energy level occupied by electron values are all positive integers (1,2,3,…) As n increases, the electron’s energy and its average distance from the nucleus increase multiple electrons are in each level so have the same n value the total number of orbitals in a level is equal to n2

1st Quantum Number Energy

2nd Quantum Number Angular Momentum Quantum Number: l indicates the shape of the orbital (sublevel) for a certain energy level, the number of possible shapes is equal to n the possible values of l are 0 and all positive integers less than or equal to n-1 each atomic orbital is designated by the principal quantum number followed by the letter of the sublevel

2nd Quantum Number s orbitals: spherical l value of 0

2nd Quantum Number p orbitals: dumbbell-shaped l value of 1

2nd Quantum Number d orbitals: various shapes l value of 2

2nd Quantum Number f orbitals: various shapes l value of 3

2nd Quantum Number Level Sublevels Sublevels 0 1 2 3 0 1 2 0 1 0

3rd Quantum Number Magnetic Quantum Number: ml indicates the orientation of an orbital around the nucleus has values from +l  -l specifies the exact orbital that the electron is contained in each orbital holds maximum of 2 electrons

4th Quantum Number Spin Quantum Number: ms indicates the spin state of the electron only 2 possible directions only 2 possible values: +½ and -½ paired electrons must have opposite spins

Energy Level 1

Energy Level 2

Energy Level 3

Energy Level 4

Electron Configurations the arrangement of electrons in an atom each type of atom has a unique electron configuration electrons tend to assume positions that create the lowest possible energy for atom ground state electron configuration- lowest energy arrangement of electrons

Rules for Arrangements Aufbau Principle- an electron occupies the lowest-energy orbital that can receive it Beginning in the 3rd energy level, the energies of the sublevels in different energy levels begin to overlap

Rules for Arrangements Pauli Exclusion Principle- no two electrons in the same atom can have the same set of 4 quantum numbers Hund’s Rule- orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second all unpaired electrons must have the same spin

Rules for Arrangements

Writing Configurations Orbital Notation: an orbital is written as a line each orbital has a name written below it electrons are drawn as arrows (up and down) Electron Configuration Notation number of electrons in sublevel is added as a superscript http://www.cowtownproductions.com/vining/Sims/atomic_electron_configurations_s1.html

Order for Filling Sublevels

Writing Configurations Start by finding the number of electrons in the atom Identify the sublevel that the last electron added is in by looking at the location in periodic table Draw out lines for each orbital beginning with 1s and ending with the sublevel identified Add arrows individually to the orbitals until all electrons have been drawn

Silicon number of electrons: 14 last electron is in sublevel: 3p 1s 2s 2p 3s 3p Valence Electrons- the electrons in the outermost energy level

Chlorine number of electrons: 17 last electron is in sublevel: 3p 2p 3s 3p

Sodium number of electrons: 11 last electron is in sublevel: 3s 1s2 2s2 2p6 3s1 1s 2s 2p 3s

Calcium number of electrons: 20 last electron is in sublevel: 4s 1s2 2s2 2p6 3s2 3p6 4s2 1s 2s 2p 3s 3p 4s

Bromine number of electrons: 35 last electron is in sublevel: 4p 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5 1s 2s 2p 3s 3p 4s 3d 4p 1s 2s 2p 3s 3p 4s 3d 4p

Argon number of electrons: 18 last electron is in sublevel: 3p 1s2 2s2 2p6 3s2 3p6 1s 2s 2p 3s 3p

Noble Gas Notation short hand for larger atoms configuration for the last noble gas is abbreviated by the noble gas’s symbol in brackets

Electron Configuration Exceptions Copper EXPECT: [Ar] 4s2 3d9 ACTUALLY: [Ar] 4s1 3d10 Copper gains stability with a full d-sublevel.

Electron Configuration Exceptions Chromium EXPECT: [Ar] 4s2 3d4 ACTUALLY: [Ar] 4s1 3d5 Chromium gains stability with a half-full d-sublevel.

Full sublevel (s, p, d, f) Half-full sublevel Stability

Bell Ringer What do you already know about how bonds are formed? Are there different types?

Bonding Introduction to Chemical Bonding

Chemical Bonds atoms rarely exist alone when atoms are bonded together, they have less potential energy and are more stable What is potential energy? chemical bond – mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together

Ionic Bonds results from electrical attraction between large numbers of cations and anions atoms donate or accept electrons from each other

Covalent Bonds results from sharing of electron pairs between two atoms the electrons shared belong to both atoms

Covalent Bonds Polar Covalent when electrons are shared unevenly Nonpolar Covalent when electrons are shared evenly

Ionic vs. Covalent

Ionic vs. Covalent bonding usually does not fall in one category or the other, but somewhere in between type of bond depends on the elements differences in electronegativities

Ionic vs. Covalent

Polarity Polar- uneven distribution of charge Show partial charges on structure by using  (lowercase delta)

Practice Determine whether each of the following bonds will be: ionic, polar covalent, OR nonpolar covalent

Practice S and H 2.5-2.1=0.4 polar covalent S and Cs 2.5-0.7=1.8 ionic C and Cl 3.0-2.5=0.5 polar covalent

Practice Cl and Ca 3.0-1.0=2.0 ionic Cl and O 3.5-3.0=0.5 polar covalent Cl and Br 3.0-2.8 nonpolar covalent

Bell Ringer How do you determine whether a compound is molecular or ionic? Give an example of each. Write the formula for the compound made from: Mg and O Ca and Br Li and N

Covalent Bonding

Molecular Compounds molecule: neutral group of atoms held together by covalent bonds molecular compound: compound whose simplest unit is a molecule

Formulas chemical formula: tells the number of each type of atom in a compound molecular formula: tells the number of each type of atom in a molecular compound ex. H2O, Cl2, C6H12O2

Molecular Compounds diatomic molecule: a molecules containing only 2 atoms usually refers to 2 of the same atoms ex: O2, Br2, F2, etc.

Formation of Covalent Bond

Formation of Covalent Bond approaching nuclei and electron clouds are attracted to each other to create a decrease in Potential Energy (PE) two nuclei and two electron clouds repel each other creating an increase in PE

Formation of Covalent Bond a distance between the nuclei is reached in which repulsion and attraction forces are equal potential energy is at the lowest point possible at the bottom of the curve on PE graph

Covalent Bonds Bond Length distance between two bonded atoms at their lowest PE average distance since there are some vibrations measured in pm (1012 pm = 1 m) stronger the bond, shorter the bond

Covalent Bonds Bond Energy energy is released when atoms become because they have lower PE the same amount of energy must be used to break the bond and form neutral isolated atoms stronger bond, higher bond energy average since varies a small amount based on atoms in entire molecule in kJ/mol

10/28 Starter Which elements naturally exist as diatomic molecules? Remember, the 7 + 1 rule How many valence electrons do each of the halogens have? Show or describe how two bromine atoms would form a covalent bond.

Octet Rule representative elements can “fill” their outer energy level by sharing electrons in covalent bonds Octet Rule- a compound tends to form so that each atom has an octet (8) of electrons in its highest energy level by gaining, losing or sharing electrons Duet Rule- applies to H and He

Octet Rule Less than 8: Boron: 6 in outer energy level More than 8: anything in 3rd period or heavier because may use the empty d orbital ex: S, P, I

Electron Dot Diagrams a way to show electron configuration identifies the number and pairing of valence electrons to show how bonding will occur write the noble gas notation identify the number of valence identify how many are paired and how many are alone do not go by Figure 6-10

Example Nitrogen 1s2 2s2 2p3 5 valence 2 are paired 3 are alone Sulfur 1s2 2s2 2p6 3s2 3p4 6 valence 4 paired (2 pairs) 2 are alone N

Lewis Structures like dot diagrams but for entire molecules atomic symbols represent nucleus and core electrons and dots or dashes represent valence electrons unshared electrons: (lone pairs) pair of electrons not involved in bonding written around only one symbol bonding electrons: written in between 2 atoms as a dash

Types of Bonds single- sharing of one pair of electrons weakest, longest double- sharing of 2 pairs of electrons stronger and shorter triple- sharing of 3 pairs of electrons strongest and shortest multiple bonds include double and triple bonds

Drawing Lewis Structures find the number of valence electrons in each atom and add them up draw the atoms next to each other in the way they will bond add one bonding pair between each connected atoms add the rest of the electrons until all have 8 (consider exceptions to octet rule)

H H C Cl H Example 1 CH3Cl methyl chloride C: 4 x 1 = 4 H: 1 x 3 = 3 Cl: 7 x 1 = 7 total = 14 electrons carbon is central H H C Cl H duet duet duet octet octet

Example 2 NH3 ammonia N: 5 x 1 = 5 H: 1 x 3 = 3 total = 8 N is central H N H H

Example 3 N2 nitrogen gas N: 5 x 2 = 10 10 electrons N N N N

Example 4 CH2O formaldehyde C: 4 x 1 = 4 H: 1 x 2 = 2 O: 1 x 6 = 6 total = 12 C is central H C H O

Polyatomic Ions charged group of covalently bonded atoms Example: CN-

NH4+ : ammonium ion

SO42- : sulfate ion 5 x 6 = 30 total = 30 + 2 = 32 OH- : hydroxide ion 6 + 1 + 1 = 8 total S O O O O O H

Example 5 O3 ozone O: 6 x 3 = 18 two completely equal arrangements the real structure is an average of these two where each bond is sharing 3 electrons instead of 4 or 2 O O O

Resonance Structures resonance – bonding between atoms that cannot be represented in on Lewis structure show all possible structures with double-ended arrow in between to show that electrons are delocalized

Example 6 NO31- N: 5 x 1 = 5 O: 6 x 3 = 18 total = 23 + 1 = 24

Covalent Network Bonding a different type of covalent bonding not specific molecules lots of nonmetal atoms covalently bonded together in a network in all directions example: diamond silicon dioxide graphite

Bell Ringer Draw the Lewis Structure for XeF4 I3- PCl5 RnCl2

Ionic Bonding

Ionic Compounds ionic bonds do NOT form molecules chemical formulas for ionic compounds represent the simplest ratio of ion types made of anions and cations

Ionic Compounds combined so that amount of positive and negative charge is equal usually crystalline solid formula of ionic compound depends of the charges of the ions combined

Formation attractive forces: oppositely charged ions nuclei and electron clouds of adjacent ions repulsive forces: like-charged ions electrons of adjacent ions

Formation distance between the ions creates a balance between those forces ions minimize their PE by combining in an orderly arrangement called a crystal lattice

Formation specific lattice pattern created depends on: charges of ions size of ions Calcium Bromide: each Ca2+ is surrounded by 8 F- each F- is surrounded by 4 Ca2+ Sodium Chloride each Na+ is surrounded by 6 Cl- each Cl- is surrounded by 6 Na+

Lattice Energy energy released when separate gaseous ion bond to form ionic solid the larger the amount of energy released, the stronger the bond since it is released, the value is negative

Ionic vs. Molecular ionic bonds and molecular bonds are both strong ionic bonds connect all ions together molecules are more easily pulled apart because intermolecular forces are weak

Ionic vs. Molecular Molecular Compounds: low melting and boiling points many are gases at room temperature Because the intermolecular forces of the molecules are weak so they are easily separated

Ionic vs. Molecular Ionic Compounds: higher melting and boiling points all are solid at room temperature hard: Because of the strong forces, it is difficult for one layer of ions to move past another brittle: if one layer is moved, the layers come apart completely

Ionic vs. Molecular Ionic Compounds: good conductors in liquid state Because ions are free to move and carry charge poor conductor in solid state Because ions are fixed in place

Bell Ringer Why is water’s structure bent and not linear?

VSEPR Theory and Molecular Shapes

VSEPR Theory V alence S hell E lectron P air R epulsion repulsion between pairs of electrons around an atom cause them to be as far apart as possible used to predict the geometry of molecules

Molecular Shapes diatomic molecules will always be linear all other molecules can have different shapes based on the number of charge clouds around the central atom charge clouds include: bonding pairs lone pairs

2 Charge Clouds no lone pairs: linear CO2

3 Charge Clouds no lone pairs: trigonal planar CH2O 1 lone pair: bent SO2

4 Charge Clouds 1 lone pair: NH3 trigonal pyramidal no lone pairs: CH4 tetrahedral 2 lone pairs: H2O bent

5 Charge Clouds no lone pairs: trigonal bipyramidal PCl5 1 lone pair: seesaw SF4