MCAT Study Guide Chemistry Ch. 4 – Particles 2017-08-15T06:45:06+00:00

I.          4.1:  ATOMS

A.     Atomic number (Z) – the number of protons in a nucleus of an atom

B.     Atomic mass number (A) – p+ + neutrons

C.    AZX

 

II.          4.2:  ISOTOPES – same element with different # of neutrons

A.     Atomic weight – weighted average of each element’s naturally occurring isotopes

 

III.          4.3:  IONS

Same element with different number of e-s

 

IV.          4.4:  NUCLEAR STABILITY AND RADIOACTIVITY

A.     Strong nuclear force

The force that holds together protons and neutrons; very strong because it has to overcome the electrical repulsion of like charges

B.     Radioactive – the state of unstable nuclei

C.    Radioactive decay

The transformation radioactive nuclei undergo to make them more stable; 3 types:

1.     Alpha decay – emission of an α particle, which consists of 2 protons and 2 neutrons

a)     Reduces atomic weight by 4 units

b)     42α

2.     Beta decay – 3 types:

a)     β- decay – conversion of neutron into proton and electron (β- particle)

(1)   Daughter product has atomic number +1 compared to parent; mass the same

(2)   0-1β

b)     β+ decay – conversion of a proton into a neutron and positron (β+ particle)

(1)   Daughter product has atomic number -1 compared to parent; mass the same

(2)   0+1β

c)     Electron capture – a nucleus captures an e- from the closest shell and converts a proton to a neutron

(1)   Daughter product has atomic number -1 compared to parent; mass the same

3.     Gamma decay – excited nucleus “relaxes” to ground state by emitting photon of EM radiation called gamma photons (γ)

a)     These have no mass nor charge

b)     00 γ

D.    Half life – time it takes for ½ of a sample of radioactive substance to decay

1.     N = N0(1/2)t/(t1/2) → N

 

V.          4.5:  ATOMIC STRUCTURE

A.     Emission spectrum

Combination of wavelengths and energies unique to each element; “fingerprint”; related to electrons getting excited then dropping to lower level, releasing energy

1.     Ephoton = hf = hc/λ s

a)     h is Planck’s constant = 6.63 x 10-34 J*s

b)     c = speed of light

c)     λ = wavelength

d)     f = frequency

B.     Bohr Model of the Atom

1.     Bohr’s model theorized that electrons had quantized energy states and traveled in certain orbits around the nucleus

2.     These electrons could “jump” to higher orbits and “drop” to lower orbits which would release energy and produce the emission spectra

3.     En = (-2.178 x 10-18 J)/n2 → calculated energy from e- moving from one level to another

4.     ΔE = h(c/λ) → emission spectra can then be calculated if you know ΔE

 C.    Quantum Model of the Atom

1.     Better describes a multi-electron atom with e-e interaction than Bohr’s model

2.     This model assigns “addresses” to es with 4 numbers (shell, subshell, orbital, spin)

a)     n – principle quantum number; each period is assigned number in ascending order

b)     l – angular quantum number, is never larger than n (tells if in s, p, f, or d)

c)     m – orbital orientation

(1)   for p orbital, there are 3; numbers assigned are -1, 0, or 1

(2)   for d orbital, there are 5; number are -2, -1, 0, 1, 2

d)     ms – magnetic spin

(1)   Either ½ or -½

D.    The Energy Shell

1.     Energy shell in quantum model is analogous to the circular orbit of Bohr model

2.     The higher the shell number, the > the energy and the > the distance from the nucleus

E.     The Energy Subshell

1.     Subshell is composed of 1 or more orbitals (denoted s, p, d, or f); these get progressively more complex and higher in energy going from s → f

2.      Orbital – the 3D region around the nucleus where the probability is highest that the e will be found

F.     The Orbital Orientation

1.     s-subshell contains 1 orientation, p-subshell contains 3 orientations

G.    The Electron Spin

1.     Every electron has 2 possible spin states (considered intrinsic magnetism)

2.     Every orbital can accommodate 2 es, one spin-up and one spin-down

3.     If an orbital is full, its electrons are “spin-paired)

 

VI.          4.6:  ELECTRON CONFIGURATIONS

A.     3 basic rules:

1.     Aufbau principle:  es occupy the lowest energy orbitals available

2.     Hund’s rule:   es in the same subshell occupy available orbitals singly before pairing up

3.     Pauli exclusion principle:  there can be no more than 2  es in any given orbital

B.     Ex:  Oxygen  → 1s22s22p4

1.     s subshell has 1 orbital; p has 3 orbitals; d has 5 orbitals; f has 7 orbitals

C.    Diamagnetic and Paramagnetic atoms

1.     Diamagnetic – an atom that has all of its electrons spin-paired

a)     These atoms will be repelled by external magnetic field

2.     Paramagnetic – atoms whose electrons are not all spin-paired

a)     These are attracted to a magnetic field

D.    Some Anomalous Electron Configurations

1.     In a few instances, atoms can achieve a lower energy state by having a completely filled or ½ filled d-subshell (that would be either 10 or 5 es)

2.     EX:  You would think Cr =

[Ar]4s23d4, but it would prefer 5 es instead of 4, so…

3.     Cr = [Ar]4s13d5 → weird!

E.     Electron Configurations of Ions

1.     Anions put their electrons into the next available subshell

a)     EX:  F- → 1s22s22p6

2.     Isoelectric – same electron configurations (Fand Cl)

3.     Valence electrons from the highest shell are always lost first

a)     EX:  Ti = [Ar] 4s23d2  and Ti+ = [Ar] 4s13d2

F.     Excited State vs. Ground State

1.     Electrons will not be in lowest energy shells

 

VII.          4.7:  GROUPS OF THE PERIODIC TABLE AND THEIR CHARACTERISTICS

Group Name Valence-shell configuration
Group I Alkali metals ns(reducing agents)
Group II Alkaline earth metals ns2 (reducing agents)
Group VII Halogens Ns2np5 (oxidizing agents)
Group VIII Noble gases Ns2np6
The d block Transition metals
The s and p blocks Representative elements
The f block Rare earth metals

 

VIII.          4.8:  PERIODIC TRENDS

A.     Shielding

1.     Each filled shell between the nucleus and the valence electrons “shield” the valence electrons from the effects of the positively charged nucleus

2.     The result is an effective reduction in the positive elementary charge (Z → Zeff)

B.     Atomic and Ionic Radius

1.     Radius decreases from left to right across a period (no new shells are added, but the attractive force between nucleus and electrons increases)

2.     Radius increases down a group because shielding increases, decreasing the effective nuclear charge the electrons feel

C.   Ionization Energy

1.     IE1 (first ionization energy) – energy needed to remove least tightly-bound electron

2.     IE increases from left to right, decreases down a group

D.    Electron Affinity

1.     EA (electron affinity) – the energy it takes to add an electron (neg if energy is released, pos if energy is needed)

2.     Typically, more neg moving left to right (except noble gases), and more neg up a group

E.     Electronegativity

1.     Electronegativity is a measure of an atom’s ability to pull electrons to itself when it forms a covalent bond

2.     Generally, increases from left to right, decreases down a group

3.     FONClBrISH

F.     Acidity

1.     Acidity is a measure of how well a compound donates protons, accepts electrons, or lowers pH

2.     Increases from right to left and increases moving down

a)     This is because larger atoms can stabilize negative charge more easily

 

Chapter 4 Summary

  • The nucleus contains protons and neutrons; their sum corresponds to the mass number [A]
  • The number of protons corresponds to the atomic number [Z]
  • An overabundance of either protons or neutrons can result in unstable nuclei which decay via the emission of various particles
  • For nuclear decay reactions, the sum of all mass and atomic numbers in the products must equal the same sum of the numbers in the reactants
  • The rate of nuclear decay is governed by a species’ half-life
  • Electrons exist in discrete energy levels within an atom; emission spectra are obtained from energy emitted as excited electrons fall from one level to another
  • The periodic table is organized into blocks based on the architecture of electron orbitals; therefore, valence electron configurations can be determined based on an element’s location in the table
  • In their ground state, electrons occupy the lowest energy orbitals available, and occupy subshell orbitals singly before pairing
  • Atoms and ions are most stable when they have an octet of electrons in their outer shell
  • The d subshell is always backfilled: for an atom in the d block of a period n, the d subshell will have a principle number of n-1
  • A half filled (d5) or filled (d10) d subshell is exceptionally stable
  • Transition metals ionize from their valence s subshell before their d subshell
  • Atomic radious increases to the left and down then periodic table; for chard species, cations <neutral atom < anions for a given element; for isoelectronic ions, the species with more protons will have a smaller radius
  • Ionization energy, electron affinity, and electronegativity increase up and to the right on the periodic table, while acidity increases to the right and down the periodic table
  • The relative electronegativities of common atoms in decreasing order are F O N Cl Br I S H

MCAT Study Guide Chemistry - Kim Matsumoto


More MCAT Study Guide Chemistry

1.

Ch. 3 Chemistry Basics

2.

Ch. 4 Atomic Structure

3.

Ch. 5 Chemical Bonds

4.

Ch. 6 Enthalpy + Entropy

5.

Ch. 7 Calorimetry + Phase Diagrams

6.

Ch. 8 Ideal Gas Law

7.

Ch. 9 Rate Laws

8.

Ch. 10 Equilibrium

9.

Ch. 11 Acids and Bases

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