Unit 3 Test Flashcards
(151 cards)
1
Q
quantum mechanical model
A
A model that explains the behavior of absolutely small particles such as electrons and photons and explains the strange behavior of electrons
2
Q
electromagnetic radiation
A
A form of energy embodied in oscillating electric and magnetic fields aka light
3
Q
amplitude
A
The vertical height of a crest (or depth of a trough) of a wave; a measure of wave intensity and determines light’s brightness, the higher it is the brighter the color gets
4
Q
wavelength (λ)
A
the distance between adjacent crests (or any two analogous points) and is measured in units such as meters, micrometers, or nanometers, determines the color
5
Q
frequency (ν)
A
For waves, the number of cycles (or complete wavelengths) that pass through a stationary point in one second, directly proportional to the speed at which the wave is traveling. Units include cycles per second (cycle/s) and Hertz (Hz)
6
Q
formula for frequency?
A
v=c/λ, c=speed of light, λ = wavelength
7
Q
electromagnetic spectrum
A
The range of the wavelengths of all possible electromagnetic radiation
8
Q
gamma (γ) ray
A
The form of electromagnetic radiation with the shortest wavelength and highest energy. produced by the sun, other stars, and certain unstable atomic nuclei on Earth. (10^-11 meters & 10-1044 J)
9
Q
X-ray
A
Electromagnetic radiation with wavelengths slightly longer than those of gamma rays; used to image bones and internal organs ( 1 x 10^-11 - 1 x 10^-8 m & 2 x 10^-17 - 2 x 10^-14 J)
10
Q
ultraviolet (UV) radiation
A
Electromagnetic radiation with slightly smaller wavelengths than visible light aka as the component of sunlight that produces a sunburn or suntan (1 x 10^-8m & 5 x 10^-19 - 2 x 10^-17J)
11
Q
visible light
A
those frequencies of electromagnetic radiation that can be detected by the human eye (4-7 x 10 -6 m & 10-19 J)
12
Q
infrared (IR) radiation
A
Electromagnetic radiation emitted from warm objects, with wavelengths slightly larger than those of visible light ( 10^-5 m 2 x 10^-22 - 3 x 10^-19J)
13
Q
microwaves
A
Electromagnetic radiation with wavelengths slightly longer than those of infrared radiation; used for radar and in microwave ovens (10^-1m)
14
Q
radio waves
A
The form of electromagnetic radiation with the longest wavelengths and smallest energy; used to transmit the signals responsible for AM and FM radio, cellular telephone, television, and other forms of communication (10^3 m)
15
Q
What is the electromagnetic spectrum from the lowest to the highest energy?
A
radio, microwave, infrared, visible, ultraviolet, x-ray, gamma
16
Q
interference
A
The superposition of two or more waves overlapping in space, resulting in either an increase in amplitude or a decrease in amplitude
17
Q
constructive interference
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The interaction of waves from two sources that align with overlapping crests, resulting in a wave of greater amplitude
18
Q
destructive interference
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The interaction of waves from two sources that are aligned so that the crest of one overlaps the trough of the other, resulting in cancellation.
19
Q
diffraction
A
The phenomena by which a wave emerging from an aperture spreads out to form a new wave front
20
Q
what’s the equation for energy?
A
E = hv or E=hc/(λ)
h, called Planck’s constant
v, frequency
c, speed of light
λ, wavelength
21
Q
photoelectric effect
A
The observation that many metals emit electrons when light falls upon them
22
Q
order the visible color spectrum from lowest wavelength to highest
A
Violet - shortest wavelength, around 380-450 nanometers with highest frequency. …
Indigo - 420 - 440 nm.
Blue - 450 - 495 nm.
Green - 495 - 570 nm.
Yellow - 570 - 590 nm.
Orange - 590 - 620 nm.
Red - longest wavelength, at around 620 - 750 nanometers with lowest frequency.
23
Q
order the visible color spectrum from lowest frequency to highest
A
Red: 400–480 THz
Orange: 480–510 THz
Yellow: 510–530 THz
Green: 530–600 THz
Blue: 600–670 THz
Indigo: 670–700 THz
Violet: 700–750 THz
24
Q
order the visible color spectrum from lowest energy to highest
A
red (limit) 1.77
red 1.91
orange 2.06
yellow 2.14
green 2.25
cyan 2.48
blue 2.75
violet (limit) 3.10
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emission spectrum
The range of wavelengths emitted by a particular element; used to identify the element; a series of discrete lines, each corresponding to a specific wavelength (and therefore energy) of light emitted by an atom or molecule
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Bohr Model
a model of the atom where the electron travels around the nucleus in circular orbit. This model's orbits exist only at specific, fixed distances from the nucleus, and the energy of each orbit is also fixed, or quantized also known as stationary states and suggested that, although they obey the laws of classical mechanics, they also possess “a peculiar, mechanically unexplainable, stability.”
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Photons
particles of light that carry energy
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Ground State
the lowest energy state of an atom, where the electrons are in their lowest possible energy levels
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Excited State
A higher energy state of an atom, where one or more electrons have absorbed energy and moved to a higher energy level
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How do atoms emit light?
Atoms emit light when electrons transition from higher energy levels to lower energy levels, releasing energy in the form of photons
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Absorption
An electron absorbs a photon and moves to a higher energy level
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Excitation
The electron is in an excited state
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Emission
The electron returns to a lower energy level, emitting a photon in the process
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de Broglie
a fundamental concept in quantum mechanics, which proposes that all matter exhibits wave-like behavior
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de Broglie wavelength (λ)
λ = h/p or λ=h/mv
λ is the wavelength,
h is Planck’s constant (6.626×10^−34)
p is the momentum of the particle
m is mass (kg)
v is velocity (m/s)
36
Heisenberg's Uncertainty Principle
a concept that states that there is fundamental limit to the precision to the certain pairs of physical properties of a particle. The more precisely one property is measured, the less precisely the other can be controlled, determined or known. For example, the more accurately the position of a particle is measured, the less accurately its momentum can be known.
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orbital
A probability distribution map, based on the quantum-mechanical model of the atom, used to describe the likely position of an electron in an atom; also, an allowed energy state for an electron.
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Schrödinger equation
an equation that describes how the quantum state of a physical system changes with time. It is essential for understanding the behavior of particles at the atomic and subatomic levels. There are two main forms of the Schrödinger equation: the time-dependent and the time-independent equations.
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quantum number
One of four interrelated numbers that determine the shape and energy of orbitals, as specified by a solution of the Schrödinger equation
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What does n represent?
Principal Quantum Number (n): Indicates the main energy level or shell. Can take positive integer values (n=1,2,3,…).
Higher n values correspond to orbitals that are farther from the nucleus and have higher energy.
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What does l represent?
Angular Momentum Quantum Number (l): Defines the shape of the orbital. Can take integer values from 0 to n−1.
Each value of l corresponds to a different type of orbital: l=0 (s orbital), l=1 (p orbital), l=2 (d orbital), l=3 (f orbital), and so on.
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What does ml represent?
Magnetic Quantum Number (m): Describes the orientation of the orbital in space. Can take integer values from −l to +l, including zero.
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What does ms represent?
Spin Quantum Number (ms): Describes the spin of the electron within the orbital in space, can take the values of + or - 1/2
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What are the different types of orbitals?
s, p, d, f
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what the s orbital?
l = 0, spherical shape, only one orientation (ml=0)
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what is the p orbital?
l = 1; dumbbell shape, three orientations (ml = -1,0 +1)
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what is the d orbital?
l = 2, more complex shapes, often described as cloverleaf or double dumbbell, five orientations (ml = -2, -1, 0, +1, +2)
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what is the f orbital?
l = 3, even more complex shapes, seven orientations (ml = -3,-2,-1,0,+1,+2,+3)
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What's the wavelength of a hydrogen transitioning from n=5 to n=2 and its color?
434 nm and purple
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What's the wavelength of a hydrogen transitioning from n=4 to n=2 and its color?
486 nm and green
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What's the wavelength of a hydrogen transitioning from n=3 to n=2 and its color?
656 nm and red
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principal level (shell)
The group of orbitals with the same value of n
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sublevel (subshell)
Those orbitals in the same principal level with the same value of n and l
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the number of sublevels in any level is equal to
n, the principal quantum number
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the number of orbitals in any sublevel is equal to
2l+1
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the number of orbitals in a level is equal to
n^2
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the number of electrons
2n^2
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how to find the change in energy that occurs in a hydrogen atom when an electron changes energy levels?
change in E = E final - Einitial
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how do i find the wavelength when an electron is transitioning between two energy levels?
wavelength = hc/E
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aufbau principle
The principle that indicates the pattern of orbital filling in an atom
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Pauli exclusion principle
The principle that no two electrons in an atom can have the same four numbers
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Hund’s rule
The principle stating that when electrons fill degenerate orbitals, they first fill them singly with parallel spins
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Coulomb’s law
A scientific law stating that the potential energy between two charged particles is proportional to the product of the charges divided by the distance that separates the charges.
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For like charges, the potential energy (E) is ---- and -------- as the particles get farther apart (as r increases). Since systems tend toward lower potential energy, like charges repel each other (in much the same way that like poles of two magnets repel each other).
positive; decreases
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For opposite charges, the potential energy is ----- and becomes more ------ as the particles get closer together (as r decreases). Therefore, opposite charges (like opposite poles on a magnet) attract each other.
negative; negative
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The magnitude of the interaction between charged particles --- as the charges of the particles increase. Consequently, an electron with a charge of 1− is more strongly attracted to a nucleus with a charge of 2+ than it is to a nucleus with a charge of 1+.
increases
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shielding
The effect on an electron of repulsion by electrons in lower-energy orbitals that screen it from the full effects of nuclear charge
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core electrons
Those electrons in a complete principal energy level and those in complete d and f sublevels
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valence electrons
Those electrons that are important in chemical bonding. For main-group elements, the valence electrons are those in the outermost principal energy leve
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isoelectronic ions
ions (or atoms) that have the same number of electrons but different nuclear charges (different numbers of protons)
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What does quantum mechanical model explain about Atomic Radius?
Generally decreases across a period due to increasing nuclear charge, pulling electrons closer, and increases down a group as electrons are added to higher energy levels
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What does quantum mechanical model explain about ionization energy?
Increases across a period due to stronger attraction between the nucleus and electrons, making it harder to remove an electron, and decreases down a group as the outer electrons are further from the nucleus
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What does quantum mechanical model explain about Electron Affinity?
Becomes more negative across a period as atoms more readily accept electrons to achieve a stable configuration, and varies down a group depending on the atomic structure
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What does the periodic table predict the metallic character of elements?
As you move from left to right across a period, the number of electrons in the outer shell increases. Elements tend to gain electrons to achieve a stable configuration (non-metallic behavior). The effective nuclear charge also increases, making it more difficult for the atoms to lose electrons. This results in a decrease in metallic character.
As you move down a group, the outer electrons are in higher energy levels further from the nucleus. The increased distance and the shielding effect of inner electrons reduce the effective nuclear charge felt by the outermost electrons, making it easier for the atoms to lose electrons. This results in an increase in metallic character.
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What does the periodic table predict the ionization energy?
Increasing Ionization Energy: As you move across a period, the number of protons in the nucleus increases, leading to a higher Z_eff. Despite the increase in electron-electron repulsion, the stronger nuclear attraction more significantly affects the outer electrons, making them harder to remove
Decreasing Ionization Energy: As you move down a group, electrons are added to higher principal energy levels (higher
𝑛
n), which are farther from the nucleus. The increased distance and additional shielding by inner electrons reduce Z_eff for the outermost electrons, making them easier to remove.
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Effective nuclear charge
the net positive charge experienced by an electron in a multi-electron atom
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Atomic and ionic radii
describe the sizes of atoms and ions
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the radii of ions with the same electron configurations decrease with an increase in ___
nuclear charge
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ionization energy
the energy required to remove an electron from an atom or ion in its gaseous state
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electron affinity
the energy change that occurs when an electron is added to a neutral atom in the gaseous state to form a negative ion. t is a measure of an atom's ability to accept an electron and form an anion
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Metallic character
properties include the tendency to lose electrons and form positive ions (cations), high electrical and thermal conductivity, malleability, ductility, luster, and a solid state at room temperature (with the exception of mercury)
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Ionic Bonds
Formed by the transfer of electrons between a metal and a nonmetal, resulting in the formation of cations and anions
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Covalent Bonds
Formed by the sharing of electrons between nonmetal atoms
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Metallic Bonds
Formed by the delocalization of electrons among metal atoms
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Ionic Compounds
Have high melting points, are hard and brittle, conduct electricity when melted or dissolved, and are often soluble in water
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Covalent Compounds
Have lower melting points, can be gases, liquids, or solids, generally do not conduct electricity, and have varying solubility in water
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Lattice Energy
The energy released when an ionic compound forms from its gaseous ions, indicating the strength and stability of the ionic bonds
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Coulomb's law
the force between two charged objects is inversely proportional to the square of the distance between them, can be used to estimate the trends in lattice energies
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Lattice energy .... as the charge of the ions increases
increases
90
Lattice energy ... as the distance between the ions increases.
decreases
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electronegativity
a measure of an atom's tendency to attract shared electrons when forming a chemical bond
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On the periodic table, electronegativity generally ... from left to right across a period and ... from top to bottom within a group
increases; decreases
93
What is the electronegativity range for ionic compounds?
Ionic: If the difference is greater than 1.7
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What is the electronegativity range for polar covalent compounds?
Polar covalent: If the difference is between 0.4 and 1.7
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dipole moment
the product of the magnitude of the charge and the distance between the centers of the positive and negative charges
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... means there is unequal sharing of electrons, while ... means equal sharing
polar, nonpolar
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formal charge
subtract the number of non-bonding electrons and half the number of bonded electrons from the number of its valence electrons
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resonance
a way to describe the bonding in certain molecules or ions by combining multiple contributing structures into a resonance hybrid
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resonance contributors
Lewis diagrams that represent different distributions of electrons for molecules that can have resonance
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resonance hybrid
the average of these structures and is considered the accurate structure for the molecule or ion
101
What are the exceptions to the octet rule?
Incomplete Octet: Some elements are stable with fewer than eight electrons in their valence shell. Common examples include:Boron (B), Aluminum (Al)
Expanded Octet: Elements in period 3 and beyond can have more than eight electrons in their valence shell due to the availability of d orbitals.
Odd Electron Species: Molecules with an odd number of electrons cannot distribute electrons to give every atom eight electrons
102
How to predict the dipole moment and indicate the δ- and δ+ bonds?
1. Determine the Molecular Geometry
Draw the Lewis Structure: Identify the arrangement of atoms and lone pairs in the molecule.
Apply VSEPR Theory: Use the Valence Shell Electron Pair Repulsion (VSEPR) theory to predict the three-dimensional shape of the molecule.
2. Identify Polar Bonds
Electronegativity Difference: Check the electronegativity values of the atoms involved in each bond. A bond is polar if there is a significant difference in electronegativity (usually greater than 0.5).
Partial Charges: The more electronegative atom will have a partial negative charge (δ-), and the less electronegative atom will have a partial positive charge (δ+).
3. Draw Bond Dipoles
Direction of Dipoles: Draw arrows pointing from the less electronegative atom (δ+) to the more electronegative atom (δ-). The length of the arrow represents the magnitude of the dipole moment.
4. Determine the Net Dipole Moment
Vector Addition: Add the bond dipoles vectorially. If the dipoles cancel each other out, the molecule is nonpolar. If they do not cancel out, the molecule has a net dipole moment and is polar.
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bond dissociation energy (bond energy)
The energy needed to break a bond in a molecule, producing two radicals ond order = ½ (bonding electrons − antibonding electrons)
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How to find the enthalpy of the reaction?
Write the Balanced Chemical Equation: Ensure the chemical equation for the reaction is balanced.
Identify Bonds Broken and Formed:
Bonds Broken: List all the bonds in the reactants that will be broken.
Bonds Formed: List all the bonds in the products that will be formed.
Use Bond Energy Data:
Bond Dissociation Energies: Use bond dissociation energy values (usually given in kJ/mol) for each type of bond.
Calculate Total Energy for Bonds Broken:
Sum the bond energies for all bonds broken. This value will be positive because energy is required to break bonds.
Calculate Total Energy for Bonds Formed:
Sum the bond energies for all bonds formed. This value will be negative because energy is released when bonds are formed.
Calculate Enthalpy Change (ΔH):
Use the formula:ΔH=∑(Bond Energies of Bonds Broken)−∑(Bond Energies of Bonds Formed)
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how to predict bond length?
inversely proportional to bond order
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electron sea model
a theory that describes the behavior of electrons in metallic bond; explains many of the unique properties of metals, including: Electrical and thermal conductivity, Luster, Malleability, Ductility, and Crystal structure but can be easily deformed; The electrons in the outer energy levels of the metal atoms are not held tightly by the atoms and can move easily from one atom to the next. This free movement of electrons is different from ionic and covalent bonds, where electrons are restricted to fixed orbitals. The attraction between the nuclear protons and the electrons in the overlapping orbitals also makes metallic bonds very difficult to break.
107
Valence Shell Electron Pair Repulsion (VSEPR) model
used to predict the geometry of molecules based on the repulsion between electron pairs in the valence shell of the central atom
Steps to Determine the Geometry:
Determine the Lewis Structure:
Draw the Lewis structure for the molecule to identify the number of bonding pairs (BPs) and lone pairs (LPs) of electrons around the central atom.
Count Electron Pairs:
Count the total number of electron pairs (both bonding and lone pairs) around the central atom.
Determine Electron Geometry:
Use the total number of electron pairs to determine the electron geometry. This gives the arrangement of electron pairs around the central atom.
Determine Molecular Geometry:
Use the number of bonding pairs and lone pairs to determine the molecular geometry, which describes the arrangement of atoms (excluding lone pairs).
108
What's the number of electron pairs, bonding pairs, lone pairs, and bond angle for linear?
2, 2, 0, 180
109
What's the number of electron pairs, bonding pairs, lone pairs, and bond angle for trigonal planar?
3, 3, 0, 120
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What's the number of electron pairs, bonding pairs, lone pairs, and bond angle for tetrahedral?
4, 4, 0, 109.5
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What's the number of electron pairs, bonding pairs, lone pairs, and bond angle for trigonal bipyramid?
5, 5, 0, 107
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What's the number of electron pairs, bonding pairs, lone pairs, and bond angle for seesaw?
5, 4, 1, 90 to 120
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What's the number of electron pairs, bonding pairs, lone pairs, and bond angle for T-shaped?
5, 3, 2, 90 and 180
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What's the number of electron pairs, bonding pairs, lone pairs, and bond angle for trigonal bipyramid (linear)?
5, 2, 3, 180
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What's the number of electron pairs, bonding pairs, lone pairs, and bond angle for octahedral?
6, 6, 0, 90
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What's the number of electron pairs, bonding pairs, lone pairs, and bond angle for square pyramidal?
6, 5, 1, 90
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What's the number of electron pairs, bonding pairs, lone pairs, and bond angle for square planar?
6, 4, 2, 90
118
Draw a 3D linear
compare with drawings
119
Draw a 3D trigonal planar
compare with drawings
120
Draw a 3D tetrahedral
compare with drawings
121
Draw a 3D trigonal pyramidal
compare with drawings
122
Draw a 3D bent
compare with drawings
123
Draw a 3D trigonal bipyramidal
compare with drawings
124
Draw a 3D seesaw
compare with drawings
125
Draw a 3D T-Shaped
compare with drawings
126
Draw a 3D linear (bipyramidal)?
compare with drawings
127
Draw a 3D octahedral
compare with drawings
128
Draw a 3D square pyramidal
compare with drawings
129
Draw a 3D square planar
compare with drawings
130
What are the non-polar shapes?
linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral, octahedral, square planar
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What are the polar shapes?
bent, trigonal pyramidal, seesaw, t-shaped, square pyramidal
132
valence bond theory
An advanced model of chemical bonding in which electrons reside in quantum-mechanical orbitals localized on individual atoms that are a hybridized blend of standard atomic orbitals; chemical bonds result from an overlap of these orbitals
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sigma (σ) bond
The resulting bond that forms between a combination of any two s, p, or hybridized orbitals that overlap end to end.
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pi (π) bond
The bond that forms between two p orbitals that overlap side to side
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Hybridized orbitals
orbitals that result from the mixing of standard atomic orbitals (s, p, d) to form new orbitals that are degenerate (of equal energy)
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sp hybridization
geometry: linear
Bond Angle: 180°
Orbitals Involved: One s and one p orbital mix to form two sp hybrid orbitals.
137
sp^2 hybridization
Geometry: Trigonal Planar
Bond Angle: 120°
Orbitals Involved: One s and two p orbitals mix to form three sp^2 hybrid orbitals.
138
sp^3 hybridization
Geometry: Tetrahedral
Bond Angle: 109.5°
Orbitals Involved: One s and three p orbitals mix to form four sp^3 hybrid orbitals
139
sp^3d hybridization
Geometry: Trigonal Bipyramidal
Bond Angle: 90°, 120°
Orbitals Involved: One s, three p, and one d orbital mix to form five sp^3d hybrid orbitals.
140
sp^3d^2 hybridization
Geometry: Octahedral
Bond Angle: 90°
Orbitals Involved: One s, three p, and two d orbitals mix to form six sp^3d^2 hybrid orbitals.
141
how to draw a 3D structure using hybrid orbital and atomic orbital overlap?
Determine the Hybridization:
Determine the hybridization of the central atom based on the number of electron pairs (regions of electron density) around it.
Identify the Geometry:
Identify the molecular geometry based on the type of hybridization.
Draw the 3D Structure:
Sketch the molecule in three dimensions, showing the spatial arrangement of the atoms.
Show Overlapping Orbitals:
Illustrate the overlapping hybrid orbitals and atomic orbitals that form the bonds.
142
molecular orbital theory
An advanced model of chemical bonding in which electrons reside in molecular orbitals delocalized over the entire molecule. In the simplest version, the molecular orbitals are simply linear combinations of atomic orbitals
143
Bonding molecular orbitals
Lower energy, electron density between nuclei, stabilize the molecule.
144
Antibonding molecular orbitals:
Higher energy, electron density outside the internuclear region, destabilize the molecule
145
Bond order
Indicates the strength and number of bonds in a molecule.
146
bond order equation
bond order = (number of bonding electrons)-(number of antibonding electrons)/2
147
nonbonding orbital
An orbital whose electrons remain localized on an atom
148
antibonding orbital
A molecular orbital that is higher in energy than any of the atomic orbitals from which it was formed.
149
diamagnetic
molecules with all electrons paired; not attracted to a magnetic field
150
paramagnetic molecules
molecules with one or more unpaired electrons; attracted to a magnetic field
151
bonding and antibonding orbitals
when atomic orbitals combine, they form molecular orbitals that can be either bonding (higher energy) or antibonding (higher energy)
