CHAPTER 4: MODERN ATOMIC THEORY
ALEKS Homework Problems: Chapter 4 (11 Topics)
4.1 Energy: the ability to do work or produce heat
Kinetic Energy (KE): energy associated with an object’s motion
– e.g. a car moving at 55 mph has much greater KE than the same car moving at 15 mph
® Greater damage if the car crashes at 55 mph than at 15 mph
– Consider the video on PE and KE: https://www.youtube.com/watch?v=Jnj8mc04r9E
Potential Energy (PE): energy due to position or its
composition (chemical bonds)
– – A 10-lb bowling ball has higher PE when it is 10 feet
off the ground compared to 10 inches off the ground
®
® More damage on your foot from greater fall
– In terms of chemical bonds, the stronger the bond,
® more energy is required to break the bond,
® the higher the potential energy of the bond
– Note that energy can be used to break the bonds
holding atoms together, but the resulting atoms
are less stable than the bonded atoms.
® Since the atoms are usually more stable
bonded together, they tend to re-form bonds.
– Thus, energy can be used to make a system less
stable (increase its PE), but that same amount of energy is released when the reverse process
is carried out to return the system to its original (more stable) state.
Law of Conservation of Energy: energy is neither created nor destroyed but transformed
Six Forms of Energy: heat, light, chemical, electrical, mechanical, and nuclear
– Each can be converted to from one form to another.
Example: Identify at least two types of energy involved for each of the following:
a. When you turn on a lamp:
__________________________________
b. When using solar panels:
__________________________________
c. At the Springfield Power Plant in The Simpsons:
__________________________________
http://cdn.appstorm.net/iphone.appstorm.net/files/2012/08/tapped_1.jpg
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Kinetic Energy and Physical States
Solids have the lowest KE of the three physical states
– Highest attraction between particles
® particles are fixed
Liquids have slightly higher KE than solids
– Particles are still attracted to each other but can move past one another
® particles are less restricted
Gases have greatest KE compared to solids and liquids
– Attractive forces completely overcome, so particles fly freely within container
® particles are completely unrestricted
KINETIC ENERGY AND HEAT
Temperature:
A measure of the average kinetic energies of the particles in a substance
– i.e., a measure of the random motions of the particles in a substance
Heat: A measure of the total energy of the particles in a system (also called thermal energy)
Thermal energy is the kinetic energy associated with the motion of particles.
– Proportional to a substance’s temperature
– Increases with the size of a sample
Example: Consider the two beakers at the right
which both contain boiling water (at 100°C).
1. Which beaker has water molecules with
higher average kinetic energy?
(a)
(b)
neither
2. Which beaker contains water with higher thermal energy?
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(a)
(b)
neither
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Heat: Energy that is transferred from a body at a higher temperature to one at a lower
temperature ® heat always transfers from the hotter to the cooler object!
– "heat flow" means heat transfer
– View “There’s No Such Thing as Cold”: https://www.youtube.com/watch?v=Akd7MMRKDwc
Heat Transfer and Temperature
– One becomes hotter by gaining heat.
– One becomes colder by losing heat—i.e., when you “feel cold”, you are actually losing heat!
Ex. 1: Fill in the blanks to indicate how heat is transferred:
a. You burn your hand on a hot frying pan.
________________ loses heat, and ______________ gains the heat.
b. Your tongue feels cold when you eat ice cream.
________________ loses heat, and ______________ gains the heat.
Ex. 2: A small chunk of gold is heated in beaker #1, which contains boiling water. The gold
chunk is then transferred to beaker #2, which contains room-temperature water.
a. The temperature of the water in beaker #2 _____.
­
¯
stays the same
b. Fill in the blanks: _____________ loses heat, and ____________ gains the heat.
Ex. 3: Why do surfaces like a stone countertop or a metal handrail inside a building feel cold?
Are they not at room temperature like you? Explain.
Energy and Chemical and Physical Change
endothermic change: a physical or chemical change that requires energy or heat to occur
– boiling water requires energy:
H2O(l) + heat energy
– electrolysis of water requires energy:
® H2O(g)
2 H2O(l) + electrical energy ® 2 H2(g) + O2(g)
exothermic change: a physical or chemical change that releases energy or heat
®
– water condensing releases energy:
H2O(g)
– hydrogen burning releases energy:
2 H2(g) + O2(g) ® 2 H2O(g) + heat energy
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H2O(l)
+
heat energy
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For physical changes, consider whether the reactants or products have more kinetic energy.
– If KE of the reactants > KE of the products ® reactants lost KE ® exothermic process
– If KE of the products > KE of the reactants ® reactants gained KE ® endothermic process
system: that part of the universe being studied
surroundings: the rest of the universe outside the system
For chemical changes, observe if the surroundings (including you) feel hotter or colder after the
reaction has occurred.
– If the surroundings become hotter, the reaction released heat ® exothermic reaction
– If the surroundings become colder, the reaction absorbed heat ® endothermic reaction
Ex. 1: Circle the following changes that are exothermic:
freezing
vaporizing
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sublimation
melting
deposition
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Ex. 2: A student adds ammonium chloride (NH4Cl) salt to a test tube containing water and notices
that the test tube feels colder as the ammonium chloride dissolves.
This process is ___________. (Choose one)
exothermic
endothermic
Ex. 3: A student mixes solutions of an acid and a base, then notices the beaker holding the
mixture becomes hotter as the chemicals mix.
This process is ___________. (Choose one)
exothermic
endothermic
Units of Energy
calorie (cal): unit of energy used most often in the US
– amount of energy required to raise the temperature of 1 g of water by 1˚C
1 cal º 4.184 J
(Note: This is EXACT!)
– But a nutritional calorie (abbreviated Cal) is actually 1000 cal: 1 Cal = 1 kcal = 4.184 kJ
Energy and Food Values
food value: The amount of heat
released when food is burned
completely, usually reported in
Cal/g food or kJ/g food.
– Most of the energy needed by our
bodies comes from carbohydrates
and fats, and the carbohydrates
decompose in the intestines to
form glucose, C6H12O6.
– The combustion of glucose produces energy that is quickly supplied to the body:
C6H12O6(g) +
6 O2(g) ®
6 CO2(g) + 6 H2O(g) + heat energy
– The body also produces energy from proteins and fats, which can be stored because fats are
insoluble in water and produce more energy than proteins and carbohydrates.
– The energy content reported on food labels is generally determined in a variety of ways
– One method a bomb calorimeter in which food is burned, and the change in temperature
provides the total energy released.
– ChemMatters: The Science Behind Calories and Nutrition Facts Labels:
https://www.youtube.com/watch?v=G0O87gWv-Xk
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joule (J): SI unit of energy
– To recognize the size of a joule, note that
1 watt = 1
J
s
® So a 100-watt light bulb uses 100 J every second.
– Heat is also often reported in kilojoules (kJ), where 1 kJ = 1000 J
specific heat capacity:
(or specific heat)
amount of heat necessary to raise the temperature of 1 gram of any
substance by 1°C; has units of J/g°C
Ex. 1: Water’s specific heat (4.184 J/g·°C) while the specific
heats of rocks and other solids (1.3 J/g·°C for dry Earth,
0.88 J/g·°C for concrete, 0.49 J/g·°C for steel).
The same mass of each sample at the same initial
temperature was heated with the same amount of
energy. Indicate which one would have the highest final
temperature. Which would have the lowest final
temperature?
highest final temp.: ___________________ lowest final temp.: ___________________
Explain why.
Thus, because water covers most of the Earth, water can absorb a lot more energy before its
temperature starts to rise.
® Water helps to regulate temperatures on Earth within a comfortable range for humans.
® Why coastal regions do not have extreme temperatures compared to desert regions
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LIGHT: ELECTROMAGNETIC RADIATION
–
–
–
Einstein’s Big Idea: Michael Faraday video: https://www.youtube.com/watch?v=V64toYdH9hU&t=367s
BBC James Clerk Maxwell video: https://www.youtube.com/watch?v=O8OUH0pPyoI&t=35m31s
NASA’s Introduction to the Electromagnetic Spectrum: https://www.youtube.com/watch?v=cfXzwh3KadE&t=13s
Light is a form of electromagnetic radiation, a type of energy that travels through space at a
constant speed, known as the speed of light (symbol c): 2.998´108 m/s (~186,000 mi./hour)
– While light may appear instantaneous to us, it’s really a wave traveling at this finite speed.
Electromagnetic Spectrum: The continuum of radiant energy (see Fig. 4.12 on p. 131)
– The substances below are about the size of the wavelength indicated in the EM spectrum.
– e.g., an atom is about 10-10-10-9 m in size while a CD is about 10-3 m (or 1 mm) thick.
visible region: the portion of the EM spectrum that we can perceive as color
For example, a "red-hot" or "white-hot" iron bar freshly removed from a high-temperature source
has forms of energy in different parts of the EM spectrum
– red or white glow falls within the visible region, heat falls within the infrared region
The term electromagnetic comes from the theory
proposed by Scottish scientist James Clerk
Maxwell that radiant energy consists of waves
with an oscillating electric field and an
oscillating magnetic field, which are
perpendicular to one another.
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Thus, these electromagnetic waves have both a wavelength and a frequency:
wavelength (l=Greek “lambda”): distance between successive peaks
frequency (n=Greek “nu”): number of waves passing a given point in 1 s
How is energy related to wavelength and frequency?
– As the wavelength ­, the frequency ¯, and the energy ¯
– As the wavelength ¯, the frequency ­, and the energy ­
Example: Circle one for each of the examples below:
a. Which has higher frequency?
red light at 700 nm
blue light at 400 nm
b. Which has higher energy?
red light at 700 nm
blue light at 400 nm
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c. Thus, wavelength and frequency
are __________ related.
directly
inversely
not
directly
inversely
not
directly
inversely
not
d. Thus, wavelength and energy
are __________ related.
e. Thus, frequency and energy
are __________ related.
Classical Descriptions of Matter
John Dalton (1803)
– Atoms are hard, indivisible, billiard-like particles.
– Atoms have distinct masses (what distinguishes on type of atom from another).
– All atoms of an element are the same.
JJ Thomson (1890s)
– discovered charge-to-mass ratio of electrons
® atoms are divisible because the electrons are one part of atom
Ernest Rutherford (1910)
– shot positively charged alpha particles at a thin foil of gold
® discovery of the atomic nucleus
James Maxwell (1873)
– visible light consists of electromagnetic waves
How Stuff Works: Classical Gas video: https://www.youtube.com/watch?v=yXsHflXB7QM
Transition between Classical and Quantum Theory
Max Planck (1900); Blackbody Radiation: https://www.youtube.com/watch?v=AjnBGWLAoZY
– heated solids to red or white heat
– noted matter did not emit energy in continuous bursts, but in whole-number multiples of
certain well-defined quantities
® matter absorbs/emits energy in bundles = "quanta"
(single bundle of energy= "quantum")
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Albert Einstein (1905); Photoelectric Effect:
(The Photoelectric Effect Applet: http://www.kcvs.ca/site/projects/JS_files/Photoelectric_Effect/PhotoElectric.html#)
– Photoelectric Effect: Light shining on a clean metal ® emission of electrons only when
the light has a minimum threshold frequency, n0
– For n < n0 ® no electrons are emitted
– For n > n0 ® electrons are emitted, more e– emitted with greater intensity of light,
– Einstein applied Planck's quantum theory to light ® light exists as a stream of
"particles" called photons
The Bohr Model: Atoms have Orbits
What is matter? The 2400-year Search for the Atom, Part 2: https://www.youtube.com/watch?v=xazQRcSCRaY&t=204s
The color display of fireworks results from atoms absorbing energy and becoming excited.
Source: http://en.wikipedia.org/wiki/Fireworks
However, atoms in an excited state are higher in energy and unstable.
® When they return to a lower, more stable energy state, they release photons (or light
energy), sometimes in the form of visible light that we can observe as colored light.
– Different elements give off different energy, which leads to their characteristic colors (e.g.
calcium for orange, barium and copper for green, lithium for reddish-pink, etc.).
– The color depends on the arrangement of electrons within each element since elements differ
in the numbers of protons and electrons.
® Elements that emit visible colors emit unique colors (see flame tests below).
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Emission Spectra: continuous or line spectra of radiation emitted by substances
– a heated solid (e.g. the filament in an incandescent light bulb) emits light that spreads out
to give a continuous spectrum = spectrum of all wavelengths of light, like a rainbow
– A rainbow is the result of white light from the Sun being reflected by water droplets.
Hydrogen Line Spectrum
– In contrast, when a sample of hydrogen is electrified, the resulting hydrogen emission
spectrum contains only a few discrete lines:
These discrete lines correspond to specific wavelengths ® specific energies
® The hydrogen atoms’ electrons can only emit certain energies
® The energy of the electrons in the atom must also be quantized.
® Planck’s postulate that energy is quantized also applies to the electrons in an atom.
– Each element has a unique line spectrum.
® Emission spectra can be used to identify unknown elements in chemical analysis.
® The element’s line spectrum is often called its "atomic fingerprint".
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A Danish physicist named Niels Bohr used the results from the hydrogen emission spectrum to
develop a quantum model for the hydrogen atom.
–
PhET Gas Discharge Lamps applet: https://phet.colorado.edu/sims/cheerpj/dischargelamps/latest/discharge-lamps.html?simulation=discharge-lamps
Bohr Postulates: Bohr Model of the Atom
1. Energy-level Postulate
– Electrons move in discrete (quantized), circular orbits around the nucleus
– "tennis ball and stairs" analogy for electrons and energy levels
– a ball can bounce up to or drop from one stair to another, but it can never sit
halfway between two levels
– Each orbit has a specific energy associated with it, indicated as the principal energy
level or quantum number, n=1, 2, 3,...
– ground state or ground level (n = 1): lowest energy state for atom
– when the electron is in the lowest energy level in a hydrogen atom
– excited state: when the electron is in a higher energy level (n = 2,3,4,...)
2. Transitions Between Energy Levels
– When an atom absorbs energy
® the electron can jump from a lower energy level to a higher energy level.
– When an electron drops from a higher energy level to a lower energy level
® the atom releases energy, sometimes in the form of visible light.
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THE QUANTUM-MECHANICAL MODEL
Limitations of the Bohr Model ® Quantum Mechanical Model
– Unfortunately, the Bohr Model failed for all other elements that had more than one proton
and more than one electron. (The multiple electron-nuclear attractions, electron-electron
repulsions, and nuclear repulsions make other atoms much more complicated than hydrogen.)
Quantum Mechanical Model
In 1920s, a new discipline, quantum mechanics, was developed to describe the motion of
submicroscopic particles confined to tiny regions of space.
– Quantum mechanics makes no attempt to specify the position of a submicroscopic particle at a
given instant or how the particle got there
– It only gives the probability of finding submicroscopic particles (e.g. food court analogy)
® Instead we “take a snapshot” of the atom at different times and “see” where the electrons
are likely to be found (See Fig. 4.17 on p. 136).
Dual Nature of the Electron
Louis de Broglie (1924)
– If light can behave like a wave and a particle
® Matter (like an electron) can behave like a wave.
– The smaller a particle, the greater its wave properties.
– Wave properties are insignificant for large objects like a baseball.
® If we throw a baseball, we can predict where it will land given its mass, velocity, etc.
– Wave properties are significant for very small particles like an electron.
® We cannot predict the motion of subatomic particles like an electron.
– Bizarre Quantum Mechanics Explained: https://www.youtube.com/watch?v=rQJ4yX1l6to&t=9s
Werner Heisenberg (1927); Heisenberg Uncertainty Principle
– For very small particles (e.g. proton, neutrons, electrons), there is an inherent uncertainty
in the particles’ position and motion.
® It is impossible to determine both the particle’s position and its momentum.
® It is impossible to determine both the position and the momentum of an electron as it
moves around an atom’s nucleus.
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ELECTRON ORBITALS AND ELECTRON CONFIGURATIONS
Every wave has a specific frequency and energy.
® The general location occupied by an electron within an atom can be predicted.
Whereas the hydrogen atom only has energy levels that are numbered (e.g. 1, 2, 3, etc.), atoms
with more than one electron have much more complicated energy levels.
® These energy levels are divided into principal energy levels, and each principal energy level is
further subdivided into sublevels designated by different letters.
Energy Levels and Sublevels
– For all other elements (with more than 1 proton and more than 1 electron), principal energy
levels (numbered 1, 2, 3,…) are further divided into energy sublevels (s, p, d, f).
– Principal Energy Level (n=1, 2, 3,…):
– Indicates the size and energy of the orbital occupied by the electron
– As n increases, the orbital becomes larger, so the electron spends more time further
away from the nucleus.
® The further the electron is from the nucleus, the higher its energy.
Principal energy levels split into energy sublevels: s, p, d, f sublevels
principal energy level (or
shell), n: n=1,2,3,...
energy sublevels: s, p, d, and f
(or subshells)
These sublevels consist of
orbitals with specific shapes
corresponding to the probability
of finding the electron in a given
region in space.
® An electron within a given energy sublevel doesn't orbit around the nucleus.
® Instead, it has a high probability of being found within a given volume corresponding to the
orbital and its energy.
ORBITALS AND THEIR SHAPES
Erwin Schrödinger (1926)
– developed a differential equation to find the electron's wave function (y), and the square of
the wave function (y2) indicates the probability of finding the electron near a given point
– probability density for an electron is called the "electron cloud"
® “shape” of atomic orbitals
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For example, for the hydrogen 1s orbital, the size of the orbital is defined by the sphere that
contains 90% of the total electron probability.
This is the probability
map for the 1s orbital,
where the darker
regions indicate where
the electron is more
likely to be found.
This is a boundary
surface representation
of the 1s orbital,
indicating the overall
volume and shape of
the orbital occupied by
the electrons in the
orbital.
s orbitals: spherical
– The size of the orbitals increase with principal quantum number, n.
® 1s < 2s < 3s, etc.
p orbitals: dumbbell-shaped
– 3 types: px, py, pz (where x, y, and z indicates axis on which orbital aligns)
– The figures below shows the probability maps and the boundary surface
representations of the p orbitals, px, pz, and py.
d orbitals:
– 5 types: dyz, dxz, dxy , d x 2 -y 2 , dz2
– These figures show the boundary surface representations of the d orbitals.
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ELECTRON CONFIGURATIONS:
– Shorthand descriptions of the arrangement of electrons within an atom
REMEMBER the following!
– s orbitals can hold 2 electrons
– a set of p orbitals can hold 6 electrons
– a set of d orbitals can hold 10 electrons
– a set of f orbitals can hold 14 electrons
Writing Electron Configurations
1. Electrons are distributed in orbitals of increasing energy, with the lowest energy orbitals filled
first. (Consider the parking garage analogy.)
2. Once an orbital has the maximum number of electrons it can hold, it is considered “filled.”
Remaining electrons must then be placed into the next higher energy orbital, and so on.
Orbitals in order of increasing energy: (See p. 141, Fig. 4.24)
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 5d < 6p
Ex. 1
He ®
Ex. 2
C
® _____ e-
C: _______________________________________
Ex. 3
S
® _____ e-
S: _______________________________________
Ex. 4
K
® _____ e-
K: _______________________________________
Ex. 5
Fe ® _____ e-
Fe: _______________________________________
_____ e-
electron configuration for He: _____________________
These are called ground state electron configurations since they represent the most stable
form of an atom in which all of its electrons are in the lowest energy levels.
– When an atom gains energy, its electrons can be excited to higher energy levels.
® In an excited state electron configuration, some lower energy levels are not filled.
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ELECTRON CONFIGURATIONS AND THE PERIODIC TABLE
Blocks of Elements
In fact, the shape of the Periodic Table corresponds to the order of energy sublevels.
– Consider the figure below to see how electrons for each element are distributed into energy
sublevels.
Electron configurations of atoms with many electrons can become cumbersome.
® Core notation using Noble gas configurations:
– Elements in the last column of the Periodic Table are called “noble gases.”
– Since noble gases are at the end of each row in the Periodic Table, all of their electrons are
in filled orbitals.
[He]
[Ne]
[Ar]
= 1s2
= 1s2 2s2 2p6
= 1s2 2s2 2p6 3s2 3p6
– Such electrons are called “core electrons” since they are more stable (less reactive) when
they belong to completely filled orbitals.
® Noble gas electron configurations can be used to abbreviate the “core electrons” of all
elements
® Electron configurations using Noble gas abbreviations are called “core notation”
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Electron Configurations using Core Notation:
a. Electron configuration for Fe using full notation:
1s2 2s2 2p6 3s2 3p6 4s2 3d6
Electron configuration for Fe using core notation: [Ar] 4s2 3d6
b. Electron configuration for Cl using full notation: ________________________________
Electron configuration for Cl using core notation: ________________________________
c. Electron configuration for Ni using full notation: ________________________________
Electron configuration for Ni using core notation: ________________________________
d. Electron configuration for Al using core notation: ________________________________
e. Electron configuration for Sr using core notation: ________________________________
f. Electron configuration for Se using core notation: ________________________________
g. Electron configuration for I using core notation: ________________________________
Note: Be able to write electron configurations for elements #1-56.
VALENCE ELECTRONS
core electrons: innermost electrons belonging to filled electron shells
valence electrons: Electrons in the outermost shell
– Since atoms want filled electron shells to be most stable, they’ll combine with other atoms with
unfilled shells (gaining or losing e–s) to get stability.
® Valence electrons lead to chemical bonds and reactions between atoms.
® An element’s chemical properties are determined by its number of valence electrons.
The electron configurations using core notation represent the core electrons with the Noble gas,
and the remaining electrons are the valence electrons.
valence electrons
For example, consider the electron configuration for Ca: [Ar] 4s2
– In Ca, the first 18 e–s are the core electrons, and the 2 e–s in 4s are valence electrons.
Ex. 1: Use core notation to write the electron configuration for chlorine: _______________
Ex. 2: A neutral chlorine atom has _____ core electrons and _____ valence electrons.
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For Main Group (A) elements, Group # ® # of valence electrons
– Elements in Group IA: Each has 1 valence electron
– Elements in Group IIA: Each has 2 valence electrons
– Elements in Group IIIA: Each has 3 valence electrons
– Elements in Group IVA: Each has 4 valence electrons
– Elements in Group VA: Each has 5 valence electrons
– Elements in Group VIA: Each has 6 valence electrons
– Elements in Group VIIA: Each has 7 valence electrons
– Elements in Group VIIIA: Each has 8 valence electrons
Electron-Dot (or Lewis) Symbols
– Show the atom of an element with
1. Element symbol representing the nucleus and core electrons
2. Dots representing the valence e–
argon
Rules for writing Electron Dot Symbol
1. Write down the element symbol
2. Determine the number of valence electrons using the group number
3. Assume the atom has four sides, and distribute electrons with one electron per side
before pairing electrons.
Write the Lewis symbol for each of the following:
boron:
phosphorus:
oxygen:
fluorine:
Although we do not delve into the quantitative aspects of the quantum-mechanical model in this
course, calculations show that atoms and ions that have the same number of valence
electrons as the noble gases (2 valence electrons for helium and 8 valence electrons for all the
other noble gases) are very low in energy and are therefore stable.
Thus, elements tend to gain or lose electrons, so they are isoelectronic with (have the same
number of electrons as) a Noble gas to become more stable.
Ex. 1a: Indicate the number of protons and electrons for the following:
Na
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Na+
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Ex. 1b: Indicate the number of protons and electrons for the following:
S2–
S
Ex. 2: Given that metals generally lose electrons and nonmetals generally gain electrons, write
the formula for the ion formed by each of the following elements:
calcium: _______
nitrogen: _______
phosphorus: _______
oxygen: _______
chlorine: _______
magnesium: _______
barium: _______
fluorine: _______
potassium: _______
isoelectronic: has the same number of electrons
Thus, Na+ is isoelectronic with _________, and S2– is isoelectronic with _________.
Ex. 1: Circle all of the following ions that are isoelectronic with argon:
K+
Sr2+
Al3+
P3-
Ti4+
Ca2+
O2-
Mg2+
Electron Configurations of Cations and Anions
For IONS, one must account for the loss or gain of electrons:
# electrons = atomic # – (charge = change in # of valence electrons)
Or you can simply use the Periodic Table
– Find out with which element the ion is isoelectronic
– Move to the left for electrons lost or to the right for electrons gained
® write the electron configuration for that element
Example 1: Fill in the blanks for the following ions:
Ion
Isoelectronic
with what
element?
Electron Config. using
core notation
Ion
Na+
I–
P–3
Se–2
Al+3
Ti+4
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Isoelectronic
with what
element?
Electron Config. using
core notation
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PERIODIC TRENDS: Atomic Size and Metallic Character
Atomic Radius (or Size): distance from the nucleus to the outermost electrons (in nm)
Periodic Trend for Atomic Radius
– Increases down a group: More p+, n, and e–
®
bigger radius
– Decreases from left to right along a period:
– Core electrons can shield the valence electrons from the attraction towards the nucleus
– Effective nuclear charge: # of protons – # of core electrons
– Number of p+ and e– increases, but electrons going into same orbitals and can’t shield
one another from the “pull of” (attraction towards) the protons in the nucleus.
– The higher the effective nuclear charge ® smaller radius because nucleus pulling e– in
Example: Compare an Al atom with a Cl atom below:
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Trend from top to bottom ® like a snowman
Trend from left to right ® like a snowman
that fell to the right
Metallic Character: Tendency to behave like a metal rather than a nonmetal
Periodic Trend for Metallic Character:
– Decreases from left to right along a period:
Metals are concentrated on the left-hand side of P.T.; nonmetals are on the right-hand side.
– Increases down a group: Looking at groups IVA and VA, go from nonmetals (C & N) to
semimetals (Si & As) to metals (Sn & Bi).
® Same snowman trends as for atomic radius!
IONIZATION ENERGY (I.E.): Energy required to remove an electron from a neutral gaseous
atom to make it an ion.
Na(g) + ionization energy
® Na+(g) + e–
Periodic Trend for Ionization Energy
– Decreases down a group:
The bigger the atom, the farther away the electrons are from the protons in the nucleus.
® The electrons are held less tightly and are more easily removed.
– Increases from left to right along a period:
– Elements with fewer (1–3) valence electrons can more easily give up electrons to gain
noble gas configuration (stability)
– Elements with more (4–7) valence electrons can more easily gain electrons to gain
noble gas configuration (stability)
Trend from top to bottom ® like an upsidedown snowman
Trend from left to right ® like a upside-down
snowman that fell to the right
Ex. 1: Rank the following elements in terms of increasing atomic radius and ionization energy:
sulfur, fluorine, oxygen, calcium, and potassium.
smallest radius = _________ < _________ <_________ < _________ <_________ = largest radius
smallest I.E. = _________ < _________ <_________ < _________ <_________= largest I.E.
CHEM 139: Bishop Chapter 4 v0122
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