ch 6 electronic structure of atoms

Question 1

quantum theory

Answer 1

When lights are turned on, the electrons in atoms get excited to a higher energy by the electricity, and emit light when they drop back down to a lower energy

Question 2

electromagnetic radiation

Answer 2

(aka radiant energy) carries energy through space. One form is visible light. All types of ER moves at speed of light (3.00*108m/s) and have wave like characteristics

Question 3

wavelength/cycle

Answer 3

distance between two adjacent peaks/troughs of a wave

Question 4

frequency

Answer 4

number of wavelengths that pass a given point each second. measured in hertz, Hz=s^-1, or /s

Question 5

frequency-wavelenth formula

Answer 5

vlambda=c. v=nu, frequency (measured in cycles/sec=s-1= /s = Hz. lambda=lambda, wavelength (measured in meters). c= speed of light=3.00*108m/s

Long wavelength=lower frequency. Higher frequency=shorter wavelength

Question 6

electromagnetic spectrum order12

Answer 6

(shortest wavelength, highest frequency) gamma rays, x rays, ultraviolet, visible light-(400-750nm)violet, blue, green, yellow, orange, red, infrared, microwaves, radio frequency (longest wavelength, lowest frequency)

Question 7

wavelength units6

Answer 7

angstrom=10^-10m
nm=10^-9m
micrometer=10^-6(mew)
mm=10^-3
cm=10^-2
m=1m

Question 8

planck

Answer 8

QUANTIZED E
discovered: energy can be either released or absorbed by atoms only in discrete “chunks” of some minimum size. Quantum (fixed amount): smallest quantity of energy that can be emitted or absorbed as electromagnetic radiation. E of a single quantum equals Planck’s constant times frequency of radiation: E=h
h (Planck’s constant)=6.626*10-34joule-seconds (J-s). Matter is allowed to emit/absorb energy only in whole number multiples of h. 3h=three quanta of energy

Question 9

einstein

Answer 9

PHOTOELECTRIC EFFECT emission of electrons from metal surfaces on which light shines. For each metal, there is a minimum frequency of light below which no electrons are emitted
Albert Einstein explains photoelectric effect: radiant energy behaves like a stream of energy packets, each energy packed called a photon. References Planck’s quantum theory to say energy of photon=E=h. Radiant energy itself is quantized

Question 10

work function

Answer 10

A certain amount of energy (work function) is required for an electron to overcome the attractive forces that hold it in the metal. If photons hitting the metal have less energy than the work function, electrons do not escape metal surface. If photons have sufficient energy, electrons are emitted. If photons have more energy than work function, excess energy appears as kinetic energy of emitted electrons

Question 11

monochromatic

Answer 11

radiation composed of a single wavelength

Question 12

spectrum (continuous, line)3

Answer 12

Spectrum is produced when radiation from sources that produce different wavelengths is separated into its different wavelength components
Continuous spectrum:no blank spots in spectrum (ex: rainbow of colors of visible light)
Line spectrum: spectrum containing radiation of only specific wavelengths. Each wavelength represented by a colored line, which are separated by black regions (wavelengths that are absent from light)

Question 13

rydberg’s equation

Answer 13

calculated the wavelengths of all the spectral lines of Hydrogen

Question 14

Bohrs equation

Answer 14

E=(-hcRH)(1/n2)h=Planck’s constant, c=speed of light, Rh=rydberg constant, n=principal quantum number. -hcRH=-2.18*1018Jfor a H atom. This equation tells how much energy the electron will have, depending on which orbit it is in. n=1→ E is most negative. As n increases, E becomes less negative

Question 15

bohrs model postulates3

Answer 15

ELECTRONS ONLY EXIST IN DISCRETE ENERGY LEVELS (QUANTIZED). ENERGY IS INVOLVED IN MOVING AN ELECTRON BETWEEN LEVELS
Only orbits of certain radii, corresponding to certain definite energies, are permitted for the electron
An electron in a permitted orbit has a specific energy and is in an “allowed” energy state. An electron in an allowed energy state will not radiate energy and therefore will not spiral into the nucleus
Energy is emitted or absorbed by the electron only as the electron changes from one allowed energy state to another. Energy is emitted/absorbed as a photon. e=h

Question 16

principle quantum number n

Answer 16

corresponds to each orbit, Orbit closest to nucleus has n=1, radius of orbit gets larger as n increases.
Ground state of atom: lowest energy state (n=1)
Excited state: higher energy state of atom, n=2 or higher
Zero-energy state: When n gets infinitely large, E=0. Represents when electron is completely separated from nucleus. Zero-energy state is higher in energy than state with negative energies’

Question 17

moving an electron between levels

Answer 17

Energy is absorbed for an electron to move to a higher state (n increases) and emitted when the electron jumps to a lower energy state (n decreases). changeE=Ef-Ei=Ephoton=hv. changeE=hv=hc/v=(-2.18*10-18J)(1/nf2-1/ni2). If nf

Question 18

de broglie

Answer 18

MATTER WAVES
an electron when moving about the nucleus is associated with a particular wavelength: lambda=h/mv. h=planck’s constant m=mass (mass of electron=9.11*10-28g), v=velocity
The quantity mv for any object is its momentum

Question 19

heisenberg

Answer 19

UNCERTAINTY PRINCIPLE
dual nature of matter places a fundamental limitation on how precisely we can know both the location and momentum of any object at the subatomic level
Heisenberg’s principle: the uncertainty principle: it is inherently impossible for us to know simultaneously both the exact momentum of the electron and its exact location in space. The more accurately one is known, the less accurately the other is known
Uncertainty principle: changex*change(mv)h/(4pi)

Question 20

schrodinger

Answer 20

WAVE FUNCTIONS
Represented by (psi). psi^2=probability density/electron density: represents the probability that the electron will be found at that location. If psi^2is large, there is high probability of finding an electron
The wave functions in schrodinger’s equation are called orbitals. Each orbital describes a specific distribution of electron density in space, given by its probability density and each orbital has its characteristic energy and shape.

Question 21

n quantum mechanical number

Answer 21

principal quantum number, n, can have positive integral values (1,2,3…). As n increases, orbital becomes larger and the electron spends more time farther from the nucleus. As n increases, electron has higher energy and is less tightly bound to nucleus. For H atom: En=-(2.18*1018J)(1/n2)as in Bohr model
collection of orbitals with same value of n are called electron shell. All the orbitals with n=3 are in the third shell.
shell with principal quantum number n will have n subshells
total number of orbitals in a shell is n^2

Question 22

l quantum mechanical number

Answer 22

The azimuthal quantum number, l, can have values from 0 to n-1, for each value of n. Defines shape of orbital. l corresponds to letters : 0-s, 1-p, 2-d, 3-f
Orbitals with the same n andl are in the same subshell. Orbitals with n=3 and l=2 are called 3d orbitals and are in the 3d subshell

Question 23

ml quantum mechanical number

Answer 23

The magnetic quantum number, ml, can have integral values between -l and l,including 0. Describes orientation of the orbital in space.
For each value of l, there are 2l+1values of ml

Question 24

s orbital

Answer 24

spherically symmetric: electron density at a given distance from nucleus is same regardless of direction in which we proceed from nucleus

Question 25

radial probability density

Answer 25

probability that we will find the electron at a specific distance from the nucleus. When plotted as a function of r, the distance from the nucleus, you get the radial probability function curve. The probability of finding the electron rises rapidly as we move away from the nucleus, maximizing at a distance of 0.529angstroms for an H atom, then falls of rapidly

Question 26

p orbital

Answer 26

concentrated in two regions (lobes) on either side of the nucleus, separated by a node at the nucleus.
For each value of n, the three p orbitals for each shell have the same size and shape but differ in spatial orientation. The orientation is represented by the subscript px,py,pz (pyrepresents a p orbital along the y axis) As n increases, p orbitals increase in size

Question 27

d orbitals

Answer 27

5 d orbitals:
dxy, dyz, dxz look like 4 leaf clovers and line on subscript plane with lobes in between axes
dx^2-y^2 looks like 4 leaf clover on xy plane with lobes along axes
dz^2 has two lobes along z axis and doughnut in xy plane.
all orbitals have the same energy

Question 28

many electron atom vs H atom

Answer 28

In a H atom, the energy of an orbital depends only on its principal quantum number, n (all the subshells have same energy). In many-electron atoms, the electron-electron repulsions cause the different subshells to be at different energies
In a many-electron atom, for a given value of n, the energy of an orbital increases with increasing value of l. Energy increases in order of ns

Question 29

George Uhlenbeck and Samuel Goudsmit

Answer 29

ELECTRONS BEHAVE IN PAIRS proposed electron spin: each electron behaves as if it were a tiny sphere spinning on its own axis.
Spin magnetic quantum number: ms, indicates the spin of electron, only has two values: +½ or -½. Spinning charge produces magnetic field. Opposite charges attract

Question 30

pauli

Answer 30

EXCLUSION PRINCIPLE: no two electrons in an atom can have the same set of four quantum numbers n, l, ml, ms. Each orbital can hold a maximum of two electrons and they must have opposite spins

Question 31

electron configuration

Answer 31

Electron configuration: way in which electrons are distributed among the various orbitals of an atom
Most stable electron configuration=ground state, electrons in lowest possible energy states. Orbitals are filled in order of increasing energy, with no more than two electrons per orbital
Electron configurations can be written by writing symbol for subshell and superscript to indicate the number of electrons in that subshell. Li (3 electrons)=1s22s1. Orbital diagrams depict orbitals by a box and each electron by a half arrow. An arrow pointing up represents positive spin and down=negative spin.

Question 32

hund

Answer 32

OCCUPY ORBITALS SINGLY for degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized. Electrons will occupy orbitals singly to the maximum extent possible (parallel spins)

Question 33

condensed electron configurations

Answer 33

using electron configuration of nearest noble-gas element of lower atomic number (noble-gas core) to abbreviate configuration. Na:[Ne]3s^1.
inner-shell electrons are called core electrons, outer shell electrons are valence electrons and are involved in chemical bonding
Elements in the same group (column) on PT have same valence electron configuration, so that is why they share many properties

Question 34

transition metals electron configuration

Answer 34

The 4s orbital is filled before the 3d orbital (corresponds to transition metals) Mn: [Ar]4s^23d^5

Question 35

rare earth elements electron configuration

Answer 35

The 14 elements corresponding to the filling of the 4f orbitals are lanthanide elements/rare earth elements. La: [Xe]6s^25d^1, Pr: [Xe]6s^24f^3
Actinide elements complete the 5f orbitals

Question 36

Electron Configurations and the Periodic Table

Answer 36

Go from right→ left and top->bottom to know order in which orbitals are filled
The s block and p block of PT together are representative elements, sometimes called the main group elements
For representative elements we do not consider completely full d or f subshells to be among the valence electrons, and for transition elements we likewise do not consider a completely full f subshell to be among the valence electrons

Question 37

abnormalities to electron configuraion

Answer 37

Transition metals in 4th column (group 6): nd^5ns^1instead of nd^4ns^2
Transition metals in 9th column (group 11):nd^10ns^1 instead of nd^9ns^1