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Universities » Iowa State University (ISU) » CHEM - CHEMISTRY » 177 - GENERAL CHEMISTRY I » Flash Cards

Exam 3 - Flashcards

Flashcard Deck Information

Class:	CHEM 177 - GENERAL CHEMISTRY I
Subject:	CHEMISTRY
University:	Iowa State University
Term:	Spring 2010

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Thermodynamically Favored

Likely to occur on its own due to thermochemical properties.

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Spontaneous

Energy must be imparted to get the reaction started.

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Entropy

The randomness in a system or the extent to which energy is distributed or dispersed among the various motions of the molecules of the system.

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Electronic Structure

The arrangement of electrons in an atom or molecule.
The number of electrons in an atom and the distribution of the electrons around the nucleus and their energies.
6.1

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Light

Consists of regularly oscillating magnetic and electrical fields.
Visible, infrared, ultra-violet, and x-ray light.
6.1

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Electromagnetic Radiation

A form of energy that has wave characteristics and that propagates through a vacuum at the characteristic speed of 3.00 x 10^8 m/s.
Visible light, radio waves, infrared radiation, and x-rays.
6.1

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Wave

A wave is a disturbance that propagates through space and time, usually with transference of energy.
6.1

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Wavelength (λ)

The distance between two adjacent peaks.
λ = c/ν
6.1

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Frequency (ν)

The number of complete wavelengths, or cycles that pass a given point each second.
ν = c/λ
6.1

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Speed of Light

3.00 x 10^8 m/s
c = ν*λ
6.1

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Photon

Particle of light
E = h*ν
6.2

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Quantization

The energy in discrete energy packets, not continuous.
6.2

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Quantum

The smallest quantity of energy that can be emitted or absorbed as electromagnetic radiation.
6.2

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Line Spectra

When atoms of a particular element excited (heat, electrical discharge) discrete colors are emitted rather than a rainbow.
6.3

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Ground State

The lowest energy state. n=1
6.3

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Excited State

When the electron is in a higher energy orbit.
n=2 or higher.
6.3

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Bohr Model

Depicts the atom as a small, positively charged nucleus surrounded by electrons that travel in circular orbits around the nucleus.
6.3

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DeBroglie Equation

λ=h/mv
v=velocity
mv=momentum
h=Planck's constant
6.4

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Electron Density Distribution

Result of calculation showing where it is likely to find electrons.
6.5

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Principle Quantum Number (n)

1,2,3, etc.
As n increases, the orbital volume increases, electron exists further from nucleus on average, electrons less tightly bound.
shell/level
6.5

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Azimuthal Quantum Number

l
1,2,3...n-1
defines the shape of the orbital
sublevel
6.5

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Magnetic Quantum Number

ml
values between L and -L
indicates orientation of orbital in space
specific orbital
6.5

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Electron Spin

Electrons behave as if they spin on their axis-giving rise to a magnetic field.
6.5

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Spin Quantum Number

ms
1/2 or -1/2
spin up or spin down
electronic spin
6.5

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Pauli Exclusion Principle

each electron in an atom has a unique set of quantum numbers: n, l, ml, and ms.
No two electrons in an atom can have the same set of 4 quantum numbers.
6.7

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Hund's Rule

For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized.
Parallel spins
6.8

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Degenerate Orbitals

orbitals at identical energy levels.
6.8

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Valence Electrons

Electrons involved in chemical bonding
All outer shell electrons
6.8

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Core Electrons

Inner shell electrons
6.8

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Thermodynamically Favored	Likely to occur on its own due to thermochemical properties.
Spontaneous	Energy must be imparted to get the reaction started.
Entropy	The randomness in a system or the extent to which energy is distributed or dispersed among the various motions of the molecules of the system.
Electronic Structure	The arrangement of electrons in an atom or molecule. The number of electrons in an atom and the distribution of the electrons around the nucleus and their energies. 6.1

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Light	Consists of regularly oscillating magnetic and electrical fields. Visible, infrared, ultra-violet, and x-ray light. 6.1
Electromagnetic Radiation	A form of energy that has wave characteristics and that propagates through a vacuum at the characteristic speed of 3.00 x 10^8 m/s. Visible light, radio waves, infrared radiation, and x-rays. 6.1
Wave	A wave is a disturbance that propagates through space and time, usually with transference of energy. 6.1
Wavelength (λ)	The distance between two adjacent peaks. λ = c/ν 6.1

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Frequency (ν)	The number of complete wavelengths, or cycles that pass a given point each second. ν = c/λ 6.1
Speed of Light	3.00 x 10^8 m/s c = ν*λ 6.1
Photon	Particle of light E = h*ν 6.2
Quantization	The energy in discrete energy packets, not continuous. 6.2

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Quantum	The smallest quantity of energy that can be emitted or absorbed as electromagnetic radiation. 6.2
Line Spectra	When atoms of a particular element excited (heat, electrical discharge) discrete colors are emitted rather than a rainbow. 6.3
Ground State	The lowest energy state. n=1 6.3
Excited State	When the electron is in a higher energy orbit. n=2 or higher. 6.3

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Bohr Model	Depicts the atom as a small, positively charged nucleus surrounded by electrons that travel in circular orbits around the nucleus. 6.3
DeBroglie Equation	λ=h/mv v=velocity mv=momentum h=Planck's constant 6.4
Electron Density Distribution	Result of calculation showing where it is likely to find electrons. 6.5
Principle Quantum Number (n)	1,2,3, etc. As n increases, the orbital volume increases, electron exists further from nucleus on average, electrons less tightly bound. shell/level 6.5

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Azimuthal Quantum Number	l 1,2,3...n-1 defines the shape of the orbital sublevel 6.5
Magnetic Quantum Number	ml values between L and -L indicates orientation of orbital in space specific orbital 6.5
Electron Spin	Electrons behave as if they spin on their axis-giving rise to a magnetic field. 6.5
Spin Quantum Number	ms 1/2 or -1/2 spin up or spin down electronic spin 6.5

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Pauli Exclusion Principle	each electron in an atom has a unique set of quantum numbers: n, l, ml, and ms. No two electrons in an atom can have the same set of 4 quantum numbers. 6.7
Hund's Rule	For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized. Parallel spins 6.8
Degenerate Orbitals	orbitals at identical energy levels. 6.8
Valence Electrons	Electrons involved in chemical bonding All outer shell electrons 6.8

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Core Electrons

Inner shell electrons 6.8

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List View: Terms & Definitions

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	Front	Back
	Thermodynamically Favored	Likely to occur on its own due to thermochemical properties.
	Spontaneous	Energy must be imparted to get the reaction started.
	Entropy	The randomness in a system or the extent to which energy is distributed or dispersed among the various motions of the molecules of the system.
	Electronic Structure	The arrangement of electrons in an atom or molecule. The number of electrons in an atom and the distribution of the electrons around the nucleus and their energies. 6.1
	Light	Consists of regularly oscillating magnetic and electrical fields. Visible, infrared, ultra-violet, and x-ray light. 6.1
	Electromagnetic Radiation	A form of energy that has wave characteristics and that propagates through a vacuum at the characteristic speed of 3.00 x 10^8 m/s. Visible light, radio waves, infrared radiation, and x-rays. 6.1
	Wave	A wave is a disturbance that propagates through space and time, usually with transference of energy. 6.1
	Wavelength (λ)	The distance between two adjacent peaks. λ = c/ν 6.1
	Frequency (ν)	The number of complete wavelengths, or cycles that pass a given point each second. ν = c/λ 6.1
	Speed of Light	3.00 x 10^8 m/s c = ν*λ 6.1
	Photon	Particle of light E = h*ν 6.2
	Quantization	The energy in discrete energy packets, not continuous. 6.2
	Quantum	The smallest quantity of energy that can be emitted or absorbed as electromagnetic radiation. 6.2
	Line Spectra	When atoms of a particular element excited (heat, electrical discharge) discrete colors are emitted rather than a rainbow. 6.3
	Ground State	The lowest energy state. n=1 6.3
	Excited State	When the electron is in a higher energy orbit. n=2 or higher. 6.3
	Bohr Model	Depicts the atom as a small, positively charged nucleus surrounded by electrons that travel in circular orbits around the nucleus. 6.3
	DeBroglie Equation	λ=h/mv v=velocity mv=momentum h=Planck's constant 6.4
	Electron Density Distribution	Result of calculation showing where it is likely to find electrons. 6.5
	Principle Quantum Number (n)	1,2,3, etc. As n increases, the orbital volume increases, electron exists further from nucleus on average, electrons less tightly bound. shell/level 6.5
	Azimuthal Quantum Number	l 1,2,3...n-1 defines the shape of the orbital sublevel 6.5
	Magnetic Quantum Number	ml values between L and -L indicates orientation of orbital in space specific orbital 6.5
	Electron Spin	Electrons behave as if they spin on their axis-giving rise to a magnetic field. 6.5
	Spin Quantum Number	ms 1/2 or -1/2 spin up or spin down electronic spin 6.5
	Pauli Exclusion Principle	each electron in an atom has a unique set of quantum numbers: n, l, ml, and ms. No two electrons in an atom can have the same set of 4 quantum numbers. 6.7
	Hund's Rule	For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized. Parallel spins 6.8
	Degenerate Orbitals	orbitals at identical energy levels. 6.8
	Valence Electrons	Electrons involved in chemical bonding All outer shell electrons 6.8
	Core Electrons	Inner shell electrons 6.8

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