Exam 1 Flashcards
4 states of matter for molecules:
SOLID
Solid: fixed position, vibrate, can’t leave position in crystalline lattice (defined shape)
4 states of matter for molecules:
LIQUID
Liquid: glued together, can slide past each other, vibrate & rotate (shape of container)
4 states of matter for molecules:
GAS
Gas: free to move, occupy entire volume container (shape & volume of container)
4 states of matter for molecules:
PLASMA
Plasma: “ionized gas” contains electrically charged particles (lightning strikes, TV)
Allotrope
Allotropes: different forms of same element
ie. oxygen gas (O2) –> ozone gas (O3)
Pure substance:
Pure substance: constant & uniform composition
Element:
Element: molecules contain only 1 type of atom
(pure sub. cannot be broken down into simpler substances by chemical changes
O2
Compounds:
Compounds: combining different atoms
(pure sub. can be broken down into simpler substances by chemical changes)
H2O
Mixture:
Mixture: 2+ types of matter (molecules) that can be present in varying amounts and can be separated by physical changes (ie. evaporation)
Homogenous mixture:
Homogenous mixture: uniform composition, appears visually same throughout
(“solution” ie. air)
Heterogeneous mixture:
Heterogeneous mixture: composition varies from point to point
distinct clumps different molecules/substance - oil & water
Table used to classify matter:
Physical properties:
Physical properties: can be observed without changing a substance into another (reversible)
mass, volume, density, boiling pt, solubility, color, softness, something melts
Intensive physical property
Intensive: independent of amount of substance present
boiling pt, density, color
Extensive physical property
Extensive: depend on amount substance present
weight, mass, length
Chemical properties:
Chemical properties: observed when matter undergoes changes in chemical composition
flammability, corrosiveness, reactivity with acid
(hint look for terms with “reacting, changing, burning”)
Changes in matter: Physical change
Physical change: don’t change composition of substance and no new substance is formed (wax melts, magnetizing solids)
Changes in matter: chemical change
result in formation of new substance with different chemical properties (combustion, oxidation)
Law of conservation of matter:
Law of conservation of matter: there is no detectable change in total quantity of matter present when matter converts from one type to another (chemical change), of changes among solid, liquid, gaseous states (physical change)
separation of mixtures techniques
Filtration: liquid separated from a solid
Substances with diff solubility can be separated using suitable solvent (sand & salt)
Substances with diff boiling points separated using distillation or evaporation
Sublimation: direct conversation from solid → gas (ammonium chloride)
Signifiant figure rules
- All non-zero digits = significant
Zeros: - Left = not significant
- Middle = significant
- Right = significant after decimal point
Rounding number rules
Adding/subtracting: same # decimal places as the # with least decimal places
Multiplying/dividing: same # of significant figures as # with least sig figs
If digit dropped < 5 (round down) if > 5 (round up)
If digit dropped = 5 (round up or down whichever yields an even value for the retained digit)
What can be used to show a larger number of sig figs?
scientific notation
All measurements have some degree of uncertainty, not exact. What are the only EXACT numbers?
counting, definition, unit conversion (infinite # sig figs, no uncertainty)
They don’t limit # sig figs in a calculation
Accuracy vs. Precision
Accuracy: very close to the true/accepted value
Precision: similar results when repeated in the same manner
Density
density = mass / volume
At particular temp & pressure, density of substance is characteristic property → often used to identify unknown substance
Determine density of irregular object: use volume water displaced in beaker
Atom:
Atom: smallest unit of an element that can participate in chemical change (indivisible)
Element:
1 type of atom
mass is a characteristic feature that is the same for all atoms of that element
Molecule:
Molecule: 2+ atoms joined by chemical bonds (could be same atom or different)
Dalton’s atomic theory:
wrong, but laid foundation for future work
- All matter is made up of tiny particles called atoms (indivisible & indestructible)
- Atoms of an element are identical in size, mass, chemical properties
- Atoms combine to form compounds in whole number ratios
4.Atoms of element cannot change into atoms another element (only rearrange)
Which chemical reactions are not possible according to Dalton’s atomic theory?
Dalton believed: atoms of a given element retain their identities in chemical reactions
CCl4 –> CH4 (not possible)
Cathode ray tube, J.J. Thomson:
Early experiment: showed electrons small negatively charged particles inside atom –> plum pudding model
Later experiment using electric filed & magnet: measure charge to mass ratio of electron
Oil drop experiment, Millikan:
measure charge on small droplets of oil by suspending them between pair of electrically charged plates
charge of oil droplet are multiples of the electron charge : e-=1.60210-19C
Alpha-ray scattering, Rutherford:
alpha particles and gold foil → discovery of nucleus in atoms (disproved the plum pudding model → let to development of modern atomic model)
Atoms much larger than nuclei & mostly empty space inside atom occupied by electrons
Nucleus has protons & neutrons (which are much much heavier than electrons) → nucleus accounts for most of an atom’s mass but very little of its size
Useful in determining the nuclear charge of the atom b/c revealed most of atom’s mass & + charge concentrated in nucleus
In a neutral atom, where does most of the mass come from?
mass atom comes from protons & neutrons, mass electrons is negligible
Rutherford’s model of atom:
- All + charge & mass concentrated inside nucleus (tiny region)
- Negatively charged particles revolve around nucleus in circular path
- Electrostatic force attraction between proton & electrons holds atom together
What were limitations to the Rutherford model of atom?
Failed to explain:
1. Stability of an atom
What did Niels Bohr study?
electromagnetic radiation
Atomic number (Z):
protons in nucleus (found in periodic table)
Neutral atom: electrons = protons
Mass number (A):
A = protons + neutrons
Nuclear symbol:
represents nucleus of an isotope
Mole:
number of atoms/molecules in a bulk sample of matter
Molar mass:
Molar mass: mass in grams of 1 mole of that substance [gmol]
Can the number of protons and neutrons in the nucleus of an atom vary?
Number or protons in the nucleus defines the element therefore is the same for all atoms of an element. However, the number of neutrons in an atom can vary → isotopes.
Isotopes:
same #protons, different #neutrons
atoms with same atomic number (Z) but different mass number (A)
carbon 12, 13, 14
Isotopes have same chemical but different physical properties (due to differences in mass)
Percentage abundance of isotopes
Atomic mass of isotopes
weighted average of isotopic masses of all the naturally occurring isotopes of an element (decimal value)
Electromagnetic (EM) radiation:
oscillating electric & magnetic field perpendicular to each other & direction propagation (ie. visible light from sun, microwaves, x-rays)
Characteristics of EM radiation: wavelength (λ)
distance between 2 consecutive peaks/troughs
Characteristics of EM radiation: frequency (ν)
cycles pass through given point / second
1/sec = Hz
Amplitude (A):
height of peak, corresponds to brightness/intensity
What speed do all types of EM radiation travel?
speed of light
c = 3 x 10^8 m/s
Relationship between frequency & wavelength
Electromagnetic spectrum: order decreasing λ
microwave > infrared > visible > ultraviolet > x-ray > gamma ray
MIVUXG
Memorize wavelength range of visible light:
ROYGBIV
red (largest λ) –> violest (smallest λ)
Photoelectric effect (equation)
light wave is particulate in nature, consisting of small packets of energy called photons
Photoelectric effect (experiment & findings)
Electrons can be ejected from surface of a metal when light have a frequency greater than some threshold shone on it
light with > threshold frequency, KE of emitted electrons increased linearly with frequency of light
KE of emitted electrons didn’t change as intensity light increased
electrons emitted directly proportional to intensity light
Photoelectric effect: threshold frequency (v0)
min frequency of light needed to eject electrons
Relationship between threshold frequency and work function
Photoelectric effect: max KE of emitted electrons
Graph of work function vs. frequency of light (photoelectric effect)
Continuous spectrum:
contains all wavelengths of visible light
Discontinuous spectrum:
missing/discontinuous wavelength (line spectra)
Line spectra (emission):
Heat sample of atoms, absorb energy and become excited (unstable → give off absorbed energy in EM radiation/light).
Atoms didn’t absorb (ground state).
Each type of atom has unique emission spectrum → identify atoms (spectroscopy)
Each emission line consists of a single wavelength of light (implies that light emitted by a gas consists of discrete energies)
Absorption spectrum:
Shine light on samples of atoms, atoms absorb light of unique wavelengths & become excited. Unabsorbed light comes out → pass it into a prism → photodetector.
Emission & absorption spectrums are photographic negatives of each other
Bohr’s model of atom:
combines classic & quantum physics
Stationary states, quantization of angular momentum, radius n’th orbital, energy levels
Bohr’s model of atom: Stationary states
electrons move around nucleus in circular path of fixed radius & energy called orbits (electron cannot live between orbits)
Bohr’s model of atom: Quantization of angular momentum
Bohr’s model of atom:
radius of the n’th orbital of what kind of atom?
applicable for mono electronic species (only 1 electron)
Bohr’s model of atom: energy of an electron in the n’th orbital of hydrogen like atom
Bohr’s model of atom:
electron jumps from low –> high orbit what happens?
In an emission spectrum of hydrogen, electron jumps from ni to nf. Find energy & wavelength of the emitted photon.
What is Bohr’s energy equation used to calculate?
Energy released when electron jump
What is Ryberg’s equation used to calculate?
Calculate wavelength of emitted photon when electron jumps from ni to nf
Bohr’s explanation of spectral lines:
Lyman series:
electron jumps from higher energy level → ground state (n = 1)
Balmer series:
electron jumps from higher energy level → n=2
Transition produces lowest λ : n=infinity → n=2 (greatest energy gap)
Transition produces highest λ: n=3 → n=2 (smallest energy gap)
As you move to higher n values how does the energy gap change?
energy gap for next shells is smaller as you move to higher n
wavelength of 1st line in Balmer series > first line Lyman series
What is the total number of spectral lines possible?
n (initial) - 1
Limitations of Bohr’s Atomic Theory
- Works only for monoelectronic atoms
- Didn’t provide any reason for why electrons can only revolve in orbits where angular moment intregral multiples…
- Didn’t provide accurate description of the electron’s location in the atom
Quantization of angular momentum
v: velocity of object
Experimental evidence for wave nature of matter:
Davisson & Germer: electrons (particles) have a diffraction pattern (characteristic of a wave)
Heisenberg uncertainty principle (HUP)
impossible to know accurately & simultaneously both the position & momentum of moving particle
Rules out existence orbits with specific radius of electrons (instead use probabilities to express electron’s position)
Shrodinger wave equation:
describe electron wave in 3D (Ψ: x, y, z)
Wavefunction: Ψ
no physical significance, but can be used to determine the distribution of the electron’s density with respect to the nucleus in an an atom,
Square of wave function
probability of finding an electron at a particular point
Solutions to shrodinger equation
set of possible wave functions (ψ), corresponding to a set of orbitals with unique energy (E)
Energy of electron in hydrogen atom important points:
Only dependent on principle quantum number (n)
Energy levels are quantized → can only have certain discrete energy values
Exact same equation for energy obtained using Bohr’s model
Orbital:
3D space around the nucleus where the probability of finding an electron is max
To define a particular orbital/wave function, how many quantum numbers do we need?
Principle quantum number (n):
size & energy of the orbital
Possible values: + integer
As value of n increases, size & energy of that orbital increases
All orbitals with same n value are in same shell
Total number of allowed orbitals with a given n value
n^2
Angular momentum quantum number (l)
shape of the orbital
for a given n, l is between 0 and n-1
(n>l)
Orbitals with same l value are in same subshell
Magnetic quantum number (ml):
orientation of the orbital
for a given subshell defined by l, ml can be: -l…0…l (l > ml)
For a particular subshell with defined l value, there are (2l+1) orbitals in subshell = possible values of ml
Electron spin quantum number (ms):
spin of the ELECTRON, says nothing about the ORBITAL
2 different orientations +1/2,-1/2 (↑, ↓)
Each orbital has a max of how many electrons?
2 electrons
Types of representations of orbitals
Node:
region with zero electron probability of finding electron
Equation for total # nodes, radial nodes, and angular nodes
total # nodes = n - 1
radial nodes = n - 1 - l
angular nodes = l
Possible orientations of 2p orbital? d orbital? (ml)
Degenerate orbitals:
same n → same amount of E
Different subshells have different energies. Arrange them in terms of increasing energy
s < p < d < f
Shielding effect
Electron electron repulsion decreases net force of attraction between nucleus and electron being removed. Reduces net + charge that an electron experiences from the nucleus.
Inner electrons shield outer electrons much more than electrons in the same shell shield each other
Penetration effect:
Ability of an orbital to attract an electron (distance to nucleus)
Closer electron is to nucleus, lower energy associated with orbital
Order penetration of orbitals vs orbital energy
Order penetration of orbitals: s>p>d>f
Order of orbital energy: s<p<d<f<
Electron configuration:
describes how electrons of an atom are filled into atomic orbitals
Aufbau principle:
fill electrons in lowest energy orbitals first
Ground state configuration: lowest energy config
Pauli’s exclusion principle:
in a given atom, 2 electrons cannot have the same set of 4 quantum numbers
Orbital hold max 2 electrons with opposite spins
Hund’s rule (degenerate orbitals):
when filling electrons into orbitals of equal energy, fill each orbital with a single electron, maintaining parallel spins (up), before doubling up electrons in that orbital set (down)
3 orbitals in p subshell have equal energy
5 orbitals in d subshell have equal energy
Electron config of cation:
loose ns or np electrons that were added last in Aufbau
EXCEPTION Transition elements: loose ns electrons before losing (n-1)d electrons
Abbreviated electron configuration:
use previous noble gas
Valence shell:
outermost shell (highest n value)
involved in chemical bonding (determine chemical properties)
Periodic table periods vs. groups
Horizontal rows: periods (n)
Vertical columns: groups → elements same group have same valence electron config (similar chemical properties)
Anomalous electron config:
Cr,Mo,Cu,Ag: unexpected, exceptions (cost additional energy)
Completely filled or ½ filled subshell: more stable (electron shifts to a higher energy orbital)
Electron configuration table
As we go across a period →
add proton to nucleus & electron to valence shell with each element
Atomic size/radius:
bond distance used to approx atomic radius b/c atoms don’t have a sharp boundary & extremely small
Periodic table trends
Nuclear charge (Z):
protons in nucleus (or magnitude of + charge)
Greater Z = greater attraction force between nucleus & valence electron = smaller atomic radius
Shielding
electron - electron repulsion = larger atomic radius
Electrons in same shell have poor shielding
Effective nuclear charge (Zeff):
pull exerted on an outer electron by the nucleus, taking into account electron-electron repulsion
Does removing an electron from an atom change the nucleuar charge?
Why are valence electrons the easiest to remove from an atom?
have highest energies, shielded more, and are farthest from nucleus
What is the determining factor for atomic radius?
Across a period? Group?
decreases as you move from left to right across a period (due to increasing nuclear charge)
increases as you move down a group (due to the increasing number of electron shells)
Isoelectronic series:
series of atoms/ions with same number of electrons therefore same shielding
(ie. O2-,F-,Na+)
Isoelectronic series: compare the radius between negative and positive ions
Most negative ion has largest radius, most positive ion has smallest radius
Explain the atomic radius trends across period & down group:
Ionization energy (IE)
energy required to remove an electron from a gaseous atom/ion
Always + for neutral atom (energy required)
Ionization energy (IE): exceptions
Move from 2A to 3A & 5A to 6A decreases IE instead of increase
first valence electron is being added to a p subshell, the full s subshell is able to shield the p subshell from the nucleus, making the first electron in a p subshell easy to remove.
first set of paired electrons is formed within a p subshell, there is a large amount of electron-electron repulsion within that orbital, which makes the fourth electron added to a p subshell easy to remove.
How does the 1st IE compare to 2nd and 3rd IE?
IE1 < IE2 < IE3 …
How can you identify the # valence electrons in an atom using IE?
seeing where a big peak in IE occurs
removing an electron from inner shell requires much much more energy than valence shell
Electron affinity:
how much an atom wants to gain an electron
(likelihood of a neutral atom to gain an electron)
negative: energy is released when an electron is added
positive: energy must be added to the system to produce an anion
zero: process is energetically neutral
Halogens: high negative EA (more likely to gain electron)
Noble gasses: large positive EA (add electron requires energy, unlikely to gain electron)
Ionization energy (IE) vs electron affinity (EA) of Mg+
Chemical bonds:
formed to decrease energy → increase stability
- Atoms only use valence electrons in bonding
- Energy must be added to break chemical bonds, forming chemical bonds releases energy
Ionic bond:
transfer of electrons (metal & nonmetal)
large electronegativity difference
Ionic crystal lattice:
orderly 3D arrangements of cations & anions held together by electrostatic force of attraction
Covalent bond:
Molecules:
Bond length:
Bond energy (BE):
Bonding pair:
Bond order:
SHARING electrons (nonmetal & nonmetal)
Molecules: smallest unit of covalent compound
Bond length: combined energy of both bonding atoms
Bond energy (BE): depth of the well at bond length → energy required to break bond or energy released when bond is formed
Bonding pair: shared electrons
Bond order: # electrons shared between a pair of atoms
Potential energy vs internuclear distance (covalent bond)
Non polar covalent bond:
equal sharing of electron pair (similar electronegativity)
Example: diatomic elements (Cl-Cl)
Polar covalent bond:
unequal sharing of electrons (closer to one of the atoms) partial negative charge (poles) NM-
ELECTRONEGATIVITY:
ability of an atom in a molecule to attract shared electron pair to itself (unmeasurable quantity, not absolute)
More electronegative atom has a stronger attraction for the bonding electrons (∂-), less electronegative atom (∂+)
Electronegativity values with respect to F
Greater electronegativity (∆EN) = greater polarity of bond
Naming: Type 1 binary ionic compounds
compound only contains 2 types of elements
metal nonmetal + ide
For cation, name = name of element
For anion, name = first part of elements name + suffix-ide
Types of metal cations:
Naming: Type 2 binary ionic compounds
contains metal that forms multiple types of cations
metal (charge on metal) nonmetal + ide
Chromium (I) nitride: Cr3N
Ionic compounds with polyatomic ions:
compound with +1 atom
Dissolved in water, stays together as a single entity
Naming: Binary covalent compounds
2 types of non metal atoms
prefix element 1 prefix element 2+ide
- First element in formula treated as cation (name using full element name)
- Second element treated anion (named similar to anion in ionic compound -ide)
Denote # atoms present use prefixes: mono (1), di (2), tri (3), tetra (4), pentra (5), hexa (6), hepta (7), octa (8)
NO (mononitrogen monoxide → nitrogen monoxide)
Lattice energy:
energy released when oppositely charged ions come together to form solid compound
When lattice energy is formed: what ions attract more slowly & release more energy
Smaller ions & ions with greater charges
Charges (q1, q2) have greater effect on LE than atomic radii (r)
Why do atoms with high ionization energy tend to form coavelent bonds?
so that they do not form ions (won’t easily give up electrons) and electrons are available for sharing.
