Pre-AP Chemistry Electrons in Atoms

ChemistryChemical BondingThe Development of Atomic ModelsDalton solid, indivisible mass

Thomson plum-pudding modelNegatively charged e- (raisins) stuck in positively charged proton doughNo neutrons

Rutherford electrons surrounding a dense nucleus

Bohr model elctrons arranged in symmetrical orbits around the nucluesplanetary modelElectrons in a given path have a fixed energy level

5. Quantum mechanical model modern mathematical description of the atom

Sodium atom:

Energy level region around the nucleus where the electron is likely to be moving.

An electron can jump from one level to another by absorbing energy.

Quantum the amount of energy required to move an electron from its present energy level to the next higher one

quantum leapQuantum mechanical model uses mathematical equations to describe the location and energy of electrons in an atomDeveloped by Erwin SchrodingerElectrons are not in definite pathsTheir location is described in terms of probability of being in a certain regionElectron cloud (ceiling fan)Conventionally, the border is drawn at 90% probabilityAtomic orbital region in space that an electron is likely to be in

Electrons can be described by a series of 4 quantum numbers.1. Principle quantum number (n)Describes the energy level Values of 1, 2, 3, 4, etc.2. Azimuthal quantum number ( l )Describes the shape of atomic orbitalsSublevelsValues of 0 to n-10 = s, 1 = p, 2 = d, 3 = fs = spherical, p = peanut shape, d&f = more complex shapes

d = daisyf = fancy

So if n = 1, then l can be 0 (s) = 1 sublevel n = 4, then l can be 0 (s), 1 (p), 2(d) , 3(f), = 4 sublevels s orbitals spherical

p orbitals dumbbell-shaped

d orbitals daisy-shaped

f orbitals fancy shapes

3. Magnetic quantum number (ml)Orientation of the orbital in spaceValues of l to +l

So s has 1 orbital p has 3 d has 5 f has 74. Spin quantum number (ms)Values of + and - Each orbital can hold 2 electrons with opposite spinsSince spinning charged objects create a magnetic field, the electrons must spin opposite directions to minimize repulsion

electron spinsEx. How many orbitals are in the following?A. 3pD. 4pB. 2sE. 3dC. 4fF. 3rd energy level

How many electrons can be in each of the above?Are these possible?nlmlms100+1/2

52-5-1/221-1+1/2Electron Configuration ways in which electrons are arranged around nuclei of atoms.Rules that govern filling of atomic orbitals:1. Aufbau principle electrons enter orbitals of lowest energy first.

2. Pauli exclusion principle An atomic orbital can describe at most two electrons. They must have opposite spins.

3. Hunds rule When electrons occupy orbital of equal energy, one electron enters each orbital until all orbitals contain one electron with parallel spins.

1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 4f

Na = 11 1s22s22p63s1

Cd =481s22s22p63s23p64s23d104p65s24d10Practice: Write the electron configuration for the following elements:Li

O

ScMore practice: Identify each of the following atoms on the basis of its electron configurations.

a) 1s22s22p6neonb) 1s22s22p63s1sodiumc) [Kr] 5s24d2zirconiumd) [Xe] 6s24f6samarium

Ground state lowest energy level for an electron. (Normal, nonexcited state)Exceptional Electron Configurations:Cr:expected: 1s22s22p63s23p64s23d4 actual: 1s22s22p63s23p64s13d5

Cu:expected: 1s22s22p63s23p64s23d9 actual: 1s22s22p63s23p64s13d10

Half-filled energy levels are more stable than other partially filled energy levels. There are other exceptions.Light and Atomic SpectraElectromagnetic radiation a series of energy waves that includes radio waves, microwaves, visible light, infrared, and ultraviolet light, X-rays, and gamma rays.

Electromagnetic Spectrum Wavelength, lcresttroughParts of a wave:Amplitude height of the wave from the origin to the crest.

Wavelength - l - distance between the crests

Frequency f (or n) the number of wave cycles to pass a given point per unit of time. The units of frequency are 1/s, s-1, or Hertz (Hz)c= lf

where c = speed of light = 3.00x108 m/s or 3.00x1010 cm/s

As l increases, f decreases.As wavelength increases, frequency decreases.

As wavelength decreases, frequency increases.

Ex. A certain wavelength of yellow light has a frequency of 2.73x1016 s-1. Calculate its wavelength. Convert to nm.C = lfl = c/fl = (3.00x108 m/s )/(2.73x1016 s-1)l = 1.10x10-8 mSpectrum series of colors produced when sunlight is separated by being passed through a prism.ROY G. BIV

Red: longest wavelength, lowest frequencyViolet: shortest wavelength, highest frequency

Atomic emission spectrum series of lines of colored light produced by passing light emitted by an excited atom through a prism. This can be used to identify the element. The atomic emission spectrum of hydrogen shows three series of lines. The lines in the UV region (Lyman series) represent electrons falling to n=1, lines in the visible region (Balmer series) represent electrons falling to n=2 and lines in the IR region (Paschen series) represent electrons falling to n=3.

Max Planck found that the energy emitted or absorbed by a body changes only in small discrete units he called quanta. He determined that the amount of radiant energy, E, absorbed or emitted by a body is proportional to the frequency of the radiation.E=hfE = energy (J)f = frequencyh = Plancks constant, 6.626x10-34 JsEinstein studied the photoelectric effect whereby light of sufficient frequency shining on a metal causes current to flow. The amplitude of the radiation was not important, the frequency was. This told him that light must be in particles, each having a given energy. Einstein proposed that electromagnetic radiation can be viewed as a stream of particles called photons:E=hfPhotoelectric EffectElectron (photons)lightmetal Energy of a photon: E = hf

Example: Calculate the energy of an individual photon of yellow light having a frequency of 2.73x1016 s-1.E=hfE = (6.626x10-34 Js)(2.73x1016 s-1)E = 1.81x10-17 J

Einsteins special theory of relativity: E=mc2Matter and energy are different forms of the same entity.

Going Further:Louis deBroglie suggested that very small particles like electrons might also display wave particles and he came up with:

deBroglies equation: l = h mv

m = mass in kgv = velocity in m/sh = Plancks constant, 6.626x10-34 Js

DeBroglies equation is used to find the wavelength of a particle. It was determined that matter behaves as through it were moving in a wave. This is important in small object such as electrons but is negligible in larger objects such as baseballs. Heavy objects have very short wavelengths.Example: Calculate the wavelength of an electron traveling at 1.24x107 m/s. The mass of an electron is 9.11x10-28g.l = h mv9.11x10-28 g 1 kg = 9.11x10-31 kg 1000gl = (6.626x10-34 Js) = (9.11x10-31 kg)(1.24x107 m/s)

l = 5.87x10-11 m

End Going Further.In the photoelectric effect, electrons (called photoelectrons) are ejected by metals (esp. alkali) when light of sufficient frequency shines on them. Red light wont work. Photoelectric cells convert light energy into electrical energy. They are used in automatically opening doors and security systems.Heisenbergs Uncertainty Principle it is impossible to determine accurately both the momentum and the position of an electron simultaneously.

We detect motion by electromagnetic radiation. This interaction disturbs electrons.

Einstein and Heisenberg!

Mendeleev- arranged elements in order of increasing atomic mass

Moseley- arranged elements in order of increasing atomic number

Periodic Law- When the elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties.

Be able to locate noble gases, representative elements, transition metals, inner transition metals.

Noble Gaseshave completely filled shells of electronssimilar electronic structuresHe1s2Ne1s2 2s2 2p6Ar1s2 2s2 2p6 3s2 3p6Kr1s2 2s2 2p6 3s2 3p6 4s2 4p6etc.65Representative Elementselements in A groups on periodic chartrepresentative - because they best represent what we know about elemental structure & periodicity

66d - Transition Elementselements in B groups on periodic chartmetalshave d electronstransition from metalsto nonmetals

67f - Transition Elementsinner transition metals

68Electron configuration using the periodic table:

Columns are called groups or families. Rows are called periods or series.

Shorthand electron configuration:

Periodic Trends in atomic size:

covalent atomic radius - half the distance between the nuclei of atoms in a homonuclear diatomic molecule (like Cl2)

Trends:

Atomic size increases going down a group because electrons are added to higher energy levels that are farther from the nucleus. It decreases going across a period because as each proton is added to the nucleus an electron is being added to the same energy level. This shell of electrons is pulled closer in towards the nucleus.

Size changes little in the transition metals because the electrons being added are core electrons.

Zeff = effective nuclear charge-actual pull of the nucleus on the valence electrons.

Zeff = Zactual - effect of e- repulsionsTrends: Increases from H to He

Decreases from He to Li because 1s electrons shield 2s electrons

Increases from Li to Be

Decreases from Be to B because 2s shield 2p

Increases from B to C to N

Decreases from N to O because of repulsion due to doubly occupied orbitals

Increases from O to F to Ne Decreases from Ne to Na because 1s,2s & 2p shield 3s

Know exceptions to Zeff trends and reasons for these.

Ionization energy (IE)

- energy required to remove the highest energy electron from a gaseous atom

Li(g) + energy Li+(g) + e-

- depends on Zeff and size

Trends: Ionization energy decreases down a a group because the valence electrons are farther from the nucleus and are thus held less tightly. Ionization energy increases across a period because atomic size decreases and valence electrons are held more tightly. Zeff increases.

1st ionization energy = energy required to remove the first electron Al + energy Al+ + e-

2nd ionization energy = energy required to remove the second electron Al+ + energy Al2+ + e-

3rd ionization energy = energy required to remove the third electron Al2+ + energy Al3+ + e-

For Na, the 1st ionization energy is fairly low but the 2nd would be high.

Know exceptions to ionization energy trends and reasons for them.

Ionic Size

Anions are larger than the atoms from which they were formed. Know why!!!

Cations are smaller than the atoms from which they were formed. Know why!!!!

Sizes of Ions Related to Positions of the Elements in the Periodic TableIsoelectronic ions- a group of ions with the same number of electronsThe one with the highest atomic number is the smallest in size (More protons pulling on the same # of electrons).

Na+, Mg2+, Ne, F-, O2-, N3- are isoelectronic. They all have 10 electrons. Mg2+ is the smallest because it has 12 protons pulling on 10 electrons. (The protons win the tug of war) N3- is the largest because it has 7 protons and 10 electrons. (The electrons win the tug of war)

Electron Affinity -energy change that occurs when an electron is added to a gaseous atom -usually exothermic

Cl(g) + e- Cl-(g)+ energy

Trends: (but many exceptions)

Electronegativity-relative tendency of an atom to attract shared electrons to itself

Trends:

FONCl (Phone Call)

Elements with high electronegativity (nonmetals ) tend to gain electrons to form anions. Elements with low electronegativities (metals) often lose electrons to form cations.

Valence electrons- electrons in the highest occupied energy level of an atom. Valence electrons are the only electrons involved in the formation of chemical bonds.

Electron dot structures for atoms: -each dot represents a valence electron p

p X s p

Examples: N N O

O Xe Xe Al

Al

Na Na I

I

Si

Si

One of the major driving forces in nature is the tendency to go to lower energy. Atoms lose, gain or share electrons to become lower in energy and thus more stable.

Metals lose electrons easily to become positively charged cations. They will usually lose their valence electrons to achieve a noble gas electron configuration.

Na Na+ + e- [Ne]3s1 [Ne]

Al Al3+ + 3e- [Ne]3s23p1 [Ne]

Some transition metals lose their highest energy level s and p electrons but still have d electrons remaining. Their electron configuration is not quite that of a noble gas but is still stable. It is called a pseudo-noble gas electron configuration. For example, zinc loses its two electrons in 4s but keeps the ten electrons in 3d.

Transition metals always lose their highest numerical energy level electrons first. Transition metals in the 4th period lose their 4s and 4p electrons before losing any from 3d. Metals in groups 3, 4, & 5 do this also.

Example:

Fe [Ar]4s23d6 Fe2+ [Ar]3d6 Fe3+ [Ar]3d5

Nonmetals tend to gain electrons to become stable and form negatively charged anions. They achieve a noble gas electron configuration.

Example: Cl + e- [ Cl ]- [Ne]3s23p5 [Ne]3s23p6

N + 3e- [ N ]3-1s22s22p3 1s22s22p6

Ionic Bonding- the attraction of oppositely charged ions (cations and anions) When the electronegativity difference between two elements is large, the elements are likely to form a compound by ionic bonding (transfer of electrons). The farther apart across the periodic table two Group A elements are, the more ionic their bonding will be.

We can use Lewis dot formulas to represent the formation of ionic compounds.

Na + Cl Na+[ Cl ]- or NaCl

Mg: N Mg2+ [ N ]3-Mg: Mg2+ or Mg3N2Mg: N Mg2+ [ N ]3-

123Properties of Ionic Compounds:They are usually crystalline solids with high melting points (>400oC)

Their molten compounds and aqueous solutions conduct electricity well because they contain mobile charged particles.

Metals

Metals form metallic solids that consist of positively charged metal cations in a sea of loosely held valence electrons. This arrangement allows metals to have their unique properties.

Metals are ductile (can be pulled into a wire) and malleable (can be hammered into a thin sheet) because the valence electrons act as grease, allowing the cations to slide past each other without colliding with each other and shattering. When ionic compounds such as NaCl are hammered, like-charged ions collide causing repulsion and the crystal shatters.

Metals can conduct electricity easily. Electricity is a flow of electrons. As electricity (electrons) enters one end of a piece of metal, an equal number of electrons exit the other end.

Alloys- solutions of solids in solids

Substitutional alloy- atoms of one metal are substituted for atoms of a similar-sized metal in a metallic crystal.Ex. brass, sterling silver, pewter

Interstitial alloy- smaller metal atoms fit into holes in the crystal structure of a metal with larger atomsEx. steel (carbon in iron)

Amalgam- alloy which contains mercury

Covalent BondingHydrogen and nonmetals of Groups 4,5,6 & 7 often become stable and gain noble gas electron configurations by sharing electrons to form covalent bonds. Atoms will usually share electrons to follow the octet rule (eight electrons, like most noble gases) or the duet rule (2 electrons, like helium).

When atoms share one pair of electrons to form a covalent bond, it is called a single covalent bond. The electrons shared between the atoms are a shared pair. A dash can be used instead of two dots to represent the shared pair. Any other electrons on the atoms are unshared pairs or lone pairs.

Ex. H2 H-H

Cl2 Cl-Cl HCl H-Cl

Atoms must sometimes share more than one pair of electrons to become stable. When two pair of electrons are shared between two atoms, it is called a double bond. If three pair are shared, it is a triple bond.

Ex. O2 N2

O = O N N

Rules for Writing Lewis Structures (electron dot structures): (Use pencil!)1.Add up the valence electrons from all the atoms. Dont worry about keeping track of which electrons come from which atoms. If you are working with an ion, you must add or subtract electrons to equal the charge.

2.Use a pair of electrons to form a bond between each pair of bound atoms.

3.Arrange the remaining electrons to satisfy the duet rule for hydrogen and the octet rule for everything else.

4.If necessary, change bonds to double or triple.

5.Remember, we cannot create or destroy electrons!

H2O 8 electrons

H-O-H

H-O-H NH3

8 electrons

H-N-H HNH4+

9-1 = 8 electrons H + H-N-H H CO2 16 electrons

O-C-O

This used 20 electrons! BAD!!!CO2 O=C=O CCl4

32 electrons CCl4

32 electrons Cl Cl-C-Cl ClCN-

9 + 1 = 10 electrons

C-N BAD!CN-

9 + 1 = 10 electrons

CN SO42-

32 electrons O 2- O S O O CO32-24 electrons

O O C O O O C O2-2-Coordinate Covalent Bond- Bond in which both electrons came from the same atom. This bond is not really any different than any other single bond.

Exceptions to the Octet Rule:

A few compounds are stable with less than an octet. They include beryllium or boron. These electron deficient compounds are very reactive.

Ex.

BF3 BeCl2

24 electrons 16 electrons

F Cl-Be-Cl F-B-FElements in the third period and below can exceed the octet rule. They can place extra electrons in empty d orbitals. Elements in the second period can not exceed the octet rule because there is no 2d orbital for the extra electrons to go into. If it is necessary to exceed the octet rule, place the extra electrons on the central atom.

Ex. PCl5 SF6

Cl Cl F F F P S Cl Cl Cl F F F

I3-

22 electrons

[ I-I-I ]-More practice: NF3 F F - N - F OF2

20 electrons

F - O - FKrF436 electrons F F- Kr - F F BeH2

4 electrons

H - Be - HSO32- 26 electrons O O - S - O2- NO3-24 electrons O O - N - O H2O214 electrons

H - O - O - HResonance occurs when more than one valid Lewis structure can be written for a molecule. The actual structure is an average of all of the resonance structures.

Ex. NO3-

O O OO - N - O O - N - O O - N - O

In nitrate, the experimental bond length is in-between that of a single bond and a double bond. It acts like a 1 1/3 bond.

Ex. Benzene, C6H6

VSEPR

Lewis structures can be used to determine the shapes of molecules. Their shapes will tell us a lot about their chemical behavior.

The valence shell electron pair repulsion (VSEPR) theory tells us that valence electrons on the central atom repel each other. They are arranged as far apart as possible around the central atom so that repulsions among them are as small as possible. When we are using VSEPR to determine molecular shape, we are really looking for regions of electron density. Double and triple bonds count the same as single bonds in determining molecular shape.

In CO2, there are only two regions of electron density (effective electron pairs) around the central atom. These regions arrange themselves as far apart as possible, making the bond angle 180o and the molecular shape linear. O = C = O

In CO32-, there are three effective electron pairs around the central atom. The bond angle will be 120o and the shape will be trigonal planar. 2- O C O O

In CH4, there are four effective electron pairs. We might expect the bond angle to be 90o. Actually, since molecules are three-dimensional, the electron pairs are 109.5o apart (further than 90o) and take a tetrahedral arrangement.

H

H C H

H

In NH3, there are four effective electron pairs. Three are shared but one is unshared. Unshared pairs of electrons take up more space than shared pair because they are pulled closer to the nucleus. The presence of the unshared pair distorts the other bond angles, making them less than 109.5o and the shape is called trigonal pyramidal. (The bond angles in ammonia are about 107o.)

H - N - H

H

In H2O, there are four effective electron pairs, also. Two are shared and two are unshared. Since the unshared pairs repel more than shared pair, the bond angle is less than 109.5o (actually 104.5o for H2O) and the shape is bent.

H O H

In PF5, there are five effective pairs, all shared. The bond angles are 90o and 120o and the shape is called trigonal bipyramidal. It is like two trigonal pyramids with their bases touching.

F F F - P - F F

In SF6, there are six effective pairs, all shared. The bond angles are 90o and the shape is called octahedral. It is like two square pyramids with their bases touching.

Molecules that exceed the octet rule and have unshared electrons can have more complex shapes such as T-shaped, see-saw, and square pyramidal.

Practice Determining Molecular Shape:

H2S

H - S - H bent,