Unit 4: Atomic structure and electrons

Copyright © Houghton Mifflin Company 1 The Rutherford atom model. A positive nucleus surrounded by electrons like our solar system. However, this model did not properly explain chemical reactivity and certain light phenomena.

In 1665 Sir Isaac Newton noticed that white (sun) light could be split into a multicolored band of light just like a rainbow. The multicolored band of light is called a color spectrum. Light

Light as a wave In the 19 th century the works of Michael Faraday and (later) James Maxwell showed that electromagnetism would produce waves which travel at the speed of light, having light waves other than those that produced the light that we could see. We now refer to this collection of different waves of electromagnetic radiation (light) as the electromagnetic spectrum (EMS)

Wavelength: distance between two consecutive peaks or troughs in a wave symbol: lambda, ; units: meters (m) Frequency: the number of waves passing a point in a given amount of time symbol: nu, ; units: 1/sec, sec -1, Hertz (Hz) Waves: wavelength and frequency

Wavelength and frequency are inversely related

c =  C = 3.00 x 10 8 m/s

Example Calculate the frequency of light with a wavelength of 5.22 x m. c = λ ν 3.00 x 10 8 m/s = 5.22 x m x ν so ν = (3.00 x 10 8 m/s) / (5.22 x m) = 5.75 x /s = 5.75 x Hz

Electromagnetic Spectrum In increasing energy, ROY G BIV

Light as particles In 1900 German scientist Max Planck found that light is given off in discrete units (quanta). He also found light energy (E) is proportional to its frequency ( ν ). The relationship is: E = h ν, where h = Planck’s constant, 6.63 x J. S.

Wave particle nature In 1905 Albert Einstein confirmed Planck’s findings and he called the quanta “photons” (packets of energy).

Example Calculate the energy of a photon of light with a frequency of 5.45 x Hz. E = h ν = 6.63 x Js x 5.45 x s -1 = 3.61 x J. You try: Calculate the energy of a photon having a wavelength of 4.5 x m.

Bohr’s observations In 1913 Niels Bohr used the observations of Planck to explain the specific lines observed in the hydrogen emission spectrum. These lines resulted from some whole number transition.

Bohr’s model Bohr suggested that a transition corresponded to an electron jumping from one possible orbit to another and emitting a photon of light energy. In Bohr’s model of the atom, the electron can only exist in these specific orbits, known as energy levels, in an atom. Normally the electron would be in its lowest available energy level, this is called its ground state.

If the atom is exposed to an energy source the electron can absorb a quantum of energy (photon) and the electron will make a quantum leap to a higher energy level. Then the electron will drop back down to a lower energy level, thereby emitting a photon of light. The energy of this photon would correspond exactly to the energy difference between the two levels.

Light that is emitted produces a unique emission spectrum.

n=1 n=2 n=3 n=4 Spectrum UV IR VisibleVisible Ground State Excited State Excited State unstable and drops back down Energy released as a photon Frequency proportional to energy drop Excited State But only as far as n = 2 this time

Emission Spectrum Animation

Line Spectra of Other Elements

Wave – particle duality Louis De Broglie (1924) proposed that ALL matter has wave and particle properties, not just electrons. Heisenberg (1927) said that because of size and speed it is impossible to know both exact position and momentum of an electron at the same time. –This is referred to as “Heisenberg Uncertainty Principle”

Quantum mechanical model Schroedinger ( ) developed the “quantum mechanical model” of the atom. He calculated the probability where to find electrons, thereby creating “electron clouds”: areas with a great chance (90 %) to find electrons. The region in space in which there is a high probability of finding an electron is now known as an “orbital”.

Orbitals Every element has discrete energy levels called principal energy levels (given with letter n). The principal levels are divided into sublevels. Sublevels contain spaces for the electron called orbitals.

Orbital types s-orbital = spherical shape, only 1 of them p-orbital = gumdrop or dumbell shape, 3 of them – one on each axis (x,y,z) d-orbital = donut shape, 5 of them f-orbital = cigar shape, 7 of them Each orbital contains a max of 2 electrons

s and p orbitals

d orbitals

Filling orbitals with electrons We have 3 general rules for “distributing” these electrons. 1.Pauli Exclusion Principal: Orbitals contain no more than two electrons. –Each electron has a spin: up (↑) or down (↓) –Two electrons must have opposite spins to occupy an orbital

2. Hund Rule: When filling orbitals, assign one electron to each orbital (of that type) before doubling up with two electrons per orbital. 3. Aufbau: Electrons fill lowest orbitals first, then proceed to higher energy levels.

First 4 energy levels

Filling orbitals Energy level Orbital type # orbitals # of types# electrons n = 1s11 s2 n = 2s, p41 s, 3 p8 n = 3s, p, d91 s, 3 p, 5 d18 n = 4s, p, d, f161 s, 3 p, 5 d, 7 f 32 Energy level = the number of orbital types Total number of orbitals in an energy level = n 2 Total number of electrons in any energy level = 2n 2

Electron configuration The electron configuration of an atom is a shorthand method of writing the location of electrons by sublevel. The sublevel is written followed by a superscript with the number of electrons in the sublevel.

Electron configuration H1s 1 He1s 2 Li1s 2 2s 1 Be1s 2 2s 2 B1s 2 2s 2 2p 1 C1s 2 2s 2 2p 2 N1s 2 2s 2 2p 3 O1s 2 2s 2 2p 4 F1s 2 2s 2 2p 5 Ne1s 2 2s 2 2p 6

Filling Diagram for Sublevels

Order of filling orbitals 1s (with 2 electrons) 2s (2), 2p (6) 3s (2), 3p (6) 4s (2), 3d (10), 4p (6) 5s (2), 4d (10), 5p (6) 6s (2), 4f (14), 5d (10), 6p (6)

Electron configuration and the Periodic Table

Practice: Give the electron configuration for: P 1s 2 2s 2 2p 6 3s 2 3p 3 Mn 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5 Br Al

Abbreviated notation When an energy level is completely filled we often use an abbreviated notation with the noble gas configuration of the last filled period representing the inner electrons. Example: Na 1s 2 2s 2 2p 6 3s 1 or [Ne]3s 1

Practice Give the abbreviated electron configuration of the following elements: S Co I

Electron configuration of Cu Cu: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 However, Cu is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 It is energetically slightly favorable for Cu to completely fill the 3d orbital, so one electron is moved from the 4s to the 3d orbital.

Electron configuration of Cr Cr shows a similar electron configuration effect as Cu. Cr is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 rather than 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 Due to the fact that a half-filled 3d orbital is energetically favorable over a filled 4s orbital.

Ion electron configuration When we write the electron configuration of a positive ion, we remove one electron for each positive charge: Na → Na + 1s 2 2s 2 2p 6 3s 1 → 1s 2 2s 2 2p 6 When we write the electron configuration of a negative ion, we add one electron for each negative charge: O → O 2- 1s 2 2s 2 2p 4 → 1s 2 2s 2 2p 6

Transition metal ions: tend to lose the s- sublevel electrons first Fe: [Ar] 4s 2 3d 6 Fe 2+ : [Ar] 3d 6 Fe 3+: [Ar] 3d 5

Valence electrons These electrons are in the outermost principal energy level of an atom: the s and p electrons beyond the noble gas core. These electrons are involved in forming bonds with other atoms Inner electrons (core electrons) are not involved in bonding

Electron dot structure Elements (except helium) have the same # of valence electrons as their group #. Electron dot structures (Lewis dot structures) are used to show valence electrons. We use one dot for each valence electron. Consider phosphorus, P, which has 5 valence electrons. Here is the method for writing the electron dot formula.

Periodic Trends

What are periodic trends? Also called “atomic trends” – take place at the atomic level Trends are general patterns or tendencies – they are general not definite – there are exceptions When looking at trends we look for increases & decreases – across  periodic – down  group

Effect on trends 1.Nuclear Charge -The “pull” of the nucleus -Proportional to the number of protons in an atom -The greater the number of protons, the stronger the nuclear charge (“pull”) -This generally affects periodic trends

Effect on trends 2.Shielding - The electron protection from the nuclear “pull” - Shield = an energy level of electrons - We are not concerned with single electrons, only energy levels of electrons - These electrons reduce the nuclear pull - Affects group trends

Effect on trends 3.Stability - Where electron arrangement is compared to stable octet (or other special stabilities) - Determines if atom gains or loses electrons -Can be used to explain anomalies in trends

Trends in Atomic Size Increases down column –Valence shell farther from nucleus because of increased shielding Decreases across period –Left to right because of the nuclear “pull” –Adding electrons to same valence shell –Valence shell held closer because more protons in nucleus

Ionization Energy Minimum energy needed to remove a valence electron from an atom –1 mole of electrons in the gaseous state (kJ/mol) The lower the ionization energy, the easier it is to remove the electron –Metals have low ionization energies

Trends in Ionization Energy Ionization Energy decreases down the group –Valence electron farther from nucleus Ionization Energy increases across the period –Left to right –Harder to remove an electron from the atom because of the increased nuclear “pull” Exceptions: Group 3, Group 6

Ionization Energy Li + energy  Li + + e - – 1 st ionization = 520 kJ/mol Li + + energy  Li +2 + e - – 2 nd ionization = 7297 kJ/mol Li +2 + energy  Li +3 + e - – 3 rd ionization = 11,810 kJ/mol Notice, each successive ionization energy is greater than the preceding one – there is a greater “pull” between the nucleus and the electron and thus more energy is needed to break the attraction.

Examining ionization energies can help you predict what ions the element will form. – Easy to remove an electron from Group IA, but difficult to remove a second electron. So group IA metals form ions with a 1+ charge.

Electron Affinity Atoms attraction to an electron It is the energy change that accompanies the addition of an electron to a gaseous atom “Opposite” of ionization energy (Concept NOT actual trend)

Across a Period – Electron affinity increases because of increased “pull” Down a Group – Electron affinity decreases because the electrons are shielded from the pull of the nucleus Exceptions: Nitrogen Group & Noble Gases

Electronegativity The ability of an atom to attract electrons when the atom is in a compound Very similar to electron affinity Across a Period – increases because of increased pull Down a Group – decreases because of shielding F is the most electronegative element

Decrease Increase

Ionic size Cations – lose electrons (positively charged) Anions – gain electrons (negatively charged) Elements gain or lose e - to become stable – being like noble gases (filled outer sublevel) IA - +1VA - -3 IIA - +2VIA - -2 IIIA - +3VIIA - -1 IVA – shareVIIIA – 0, stable

Ionic size GOOD RULE OF THUMB – Anions are always larger than their neutral atom – Cations are always smaller than neutral atom Across a Period: Cations decrease (I-III) because of greater pull on electrons Anions decrease (V-VII) because of less pull on electrons and repulsion of the electrons Down a Group: Both cations and anions increase size

Reactivity Reactivity of metals increases to the left on the Period and down in the column –follows ease of losing an electron Reactivity of nonmetals (excluding the noble gases) increases to the right on the Period and up in the column