Structure Of Matter Flashcards
Antoine Lavoisier
Credited with the formulation of the law of conservation of matter in 1789
Law of conservation of mass
In a chemical reaction, matter is niether created nor destroyed
Key concept to the formulation of modern atomic theory
Joseph Proust
Credited with the formulation of the law of definite proportions in 1797
Law of definite proportions
All samples of a given compound have the same proportions of their constituent elements
Mass ratio
The number of each elements in the compound if the fewest elements number just one
Eg: nitrogen to carbon ratio of 4.7:1
John Dalton
Credited with the formulation of the law of multiple proportions in 1804, and then the atomic theory in 1808
(Many other works too, should probobly check wikipedia for more)
Law of multiple proportions
When two elements form two different compounds the masses of the elements can be described in a mass ratio
Atomic theory
1) each element is composed of tiny, indestructable particles called atoms
2) all atoms of a given element have the same mass and other properties that make them unique from any other element
3) atoms combine in simple, whole number ratios to form compounds
4) atoms od one element cannot change into atoms of another element, but can change the way in which they are bonded to other atoms
Atomic mass
The average mass of one mole (6.022x10^23) of each element
Given in amu (equals g/mol)
Not constant from one sample to another
Sometimes given as a range in which that mass could fall
Physical determination of atomic mass
Sum of the products of each isotope’s decimal abundance, mass of each isotope
Mass spectrometery
Method of separating atoms by their mass
Sample injected into cylinder, vaporized, ionized by electrons, accelerated through a magnetic field.
The change in their trajectory due to this magnetic force tells us their individual masses, lighter ions experience greater change
Atomic number (z)
Tells 1) number of protons 2) number of electrons (when nuetral) 3) positive charge of the nucleus Calculated as z=A-N
Neutron Count (N)
Number of neutrons in a given atom
Calculated as N=A-z
Mass number (A)
Total number of nucleons (protons and neutrons) in an atom A=z+N Often written as [Element name]-A (eg: carbon-12)
Isotopes
Atoms with the same number of protons but different neutrons
Isotope notation
A over z and then the element name
Mass number is always written over the number of protons
Natural abundance
Percentage breakdown of isotopes present in any particular sample of an element, varies by isotope
Electron energy levels
Each row of the periodic table is a new energy level for electrons
The ‘unfilled level’ being the level containing valance electrons
Electron configuration
Way of showing which electrons occupy which orbitals
Ground state= H 1s^1
Pauli exclusion principal
No two electrons can ever have the same four quantum numbers
Orbital
Any one of four areas that describe the entire distribution of where that electron might be located.
P orbital
Electrons here occupy an any space within a peanut shape arround the nucleus
D orbital
Electrons here occupy an any space within a flower shape arround the nucleus
F orbital
Electrons here occupy an any space within a 3D-flower shape arround the nucleus
Quantum numbers
Values that describe the behavior of an individual electron within an atom
S orbital
Lowest energy
Electrons here may occupy a sphere-like area around the nucleus
Principal quantum number (n)
Positive integer that represents electron energy level
Angular momentum quantum number(l)
Integer that represents the shape of the orbital
S=0
P=1
D=2
F=3
Magnetic quantum number (m sub-l)
Integer that represents orbital orientation
Exist on the interval (-l to +l)
l is the angular momentum quantum number
Spin quantum number (m sub-s)
Refers to the direction in which the electron spins
Either -1/2 or +1/2
Periodic property
A property of an atom predictable by its position on the periodic table
Electron configuration
A way of writing which electrons are in which orbital
‘Ground State’- 1s^1
Dmitry Mendeleev
Russian chemist Credited with the arrangement of the periodic table based on the work of German chemist Julius Lothar Meyer
Orbital Diagram
Way of drawing the ‘direction’ of electron ‘spin’ within electron configuration
Little arrows going up and down
Pauli Exclusion Principal
No two electrons can have the same four quantum numbers within the same atom
Coulombs law
E=q1q2/r1/(4(pi)(8.85x10^-12))
The energy between two charged particles is equal to the product of their charges divided by the distance between them over 4pi*energy constant
1) energy decreases with distance
2) opposite charges attract one another, while particles of opposite charges repel
3) particles are attracted to one another based on the magnetude of the opposing charge
Shielding
The property of an electron to repel another electron
Effective Nuclear charge
Refers to the charge acting on a particular electron
The ionic charge of that atom if the electron did not exist
Penetration
The action of an electron changing energy levels due to an attraction to the Nucleus
Aufbau Principal
Only two electrons of opposite spins are allowed in each orbital
Hunds Rule
Electrons must fill these orbitals as single electrons before they can double up
Valance Electrons
Electrons that are exchanged during chemical bonding
Core electrons
Electrons that would compose an atom of the noble nearest two but less than the mass of the atom in question
Orbital Blocks of the periodic table
A way of demonstrating on the periodic table which orbitals are used in electron configuration
G1and2 = s
Transitions = d
Metaloids/nonmetals = p
Weird rows at the bottom = f
Noble gasses
Elements of the row furthest to the right of the periodic table
Rarely bond with any other element
Used in experiments in which a reaction is undesireable
Alkali Metals
Group 1 elements, highly reactive
One valence electron
Excelent reducing agents (decrease charge of a material by contributing an electron)
Result in violent reactions
Alkali Earth Metals
Group 2 elements
Two valence electrons
Halogens
Group 7
7 valence electrons
Excelent oxideizing agents (take electrons from a material thereby increasing their charge)
Van der wals radius
The radius of an atom when it is not bonded to another atom
Covalent radius
Radius of an atom when it is bonded to another atom
Atomic radii
Lower left of the periodic table
Electronegativity
Ability for an atom to attract electrons
Opposite of metallic Character
Increases to the upper right of an atom
Metallic Character
Ability for an atom to loose an electron
Opposite of electronegativity
Increases to the lower left of the periodic table
Paramagnetism
Property of atoms to be attracted to an external magnetic field
An effect of an empty space in a low energy orbital
Usually attributed to silver
Diamagnetic
Property of atoms to be repelled by an external magnetic field
An effect of full energy orbitals
Usually attributed to zinc
Cation radius rule
Cations are nearly half the size of their nuetral atoms
Anion radius rule
Anions are much larger than their neutral atoms
Electron affinity
The energy change associated with gaining an electron in a gaseous state
Almost always negative
Increases (from negative… Approaches zero) to the upper left of the periodic table
Ionic bonds
Oppositely charges ions attract one another for an exceptionally low net charge
Electrons are tranfered in this bond
Electronegativity <.5
Covalent bonds
Two nonmetals bonding together
Form molecules
Atoms share their electrons
Electronegativity .5
Metallic bonds
Metals bond to an atom of the same element by pooling their vallence electrons (electron sea model)
No positive metal ions are attracted to the ‘sea’ of electrons
Get clarification on this
Macromolecular bonding
Smaller molecules bonded together in a polymer chain
Hydrogen bonding
Strong dipole-dipole bond between Hydrogen and either O, N, or F
Polar covalent bonds
High electronegativity difference between two atoms bonded together
Electronegativity > 1.7
Pi bond
Bond that forms betwwen two overlapping P orbitals
Peptide bond
Bond that forms between the amine end of one amino acid and the carboxylic end of another
Sigma bond
Bond resulting from an overlap of a p orbits and an s orbital
Commonly sp^2 hybridization
Valence bond theory
Advance bonding model in which electrons reside in quantum mechanicalorbits localized in individual atoms, a hybridized blend of standard atomic orbitals
A bond occurs when these orbitals overlap
Lewis theory
Simple model of of chemical bonding in which atoms are arranged to form ‘octets’ for each atom
Lewis structures
Drawings of atoms in which bonds are represented as dots to form octets
Octet rule
When bonded together, electrons must share as many electrons as it takes so that they each have ready access to 8 electrons
Lewis base
An atom, ion, or molecule that donates an electron pair
Lewis acid
An atom, ion, or molecule that accepts an electron pair
Double bond
Bond that forms when two electrons are shared between atoms
Triple bond
Bond that forms when three electrons are shared between atoms
Single bond
Bond that forms when just one electron is shared between atoms
Bonding pair
Pair of electrons shared between atoms
Bonding orbital
A molecular orbital that is lower in energy than any atomic orbitals from which it was formed
Bond order
((Number of electrons in bonding orbitals)-(Number of electrons in non-bonding orbitals))/2
Only stable when positive
Stable molecular bond
A molecule for which there are more bonding electrons than non-bonding electrons
Bond length
Average length of a bond between two particular atoms in a variety of compounds
Bond energy
The energy required to break 1 mol of the bond in a gasseous state
Chemical bond
Sharing or transfer of electrons to form stable electron configurations for bonding atoms
Molecular Orbital theory
Advance model of molecular bonding in which atomic orbitals are delocalized over the molecule as a whole, giving the molecule its own orbital
Molecular geometry
The geometric arrangement of atoms in a molecule
Valence shell electron pair repulsion (VSEPR) theory
Bonds themselves repel one another within a molecule
Electron groups
Any form of bond or even a single atom
Linear molecular geometry
Bonds are set up 180 degrees from one another
Trigonal planar geometry
Bonds are arranged 120 degrees from one another on a single plane.
Tertrahedral geometry
Bonds form a 4 sided pyramid shape within the molecule, bonds are spaced 109.5 degrees from one another
Bipyramidal geometry
Within the molecule, 3 bonds exist on one plane 120 degrees from one another, with an additional bond at either end spaced 90 degrees from the plane
Octahedral Geometry
In a molecule, four of the bonds exist on a plane while another rests at either end of and 90 degrees from this plane
Electron geometry
Geometrical arrangement of the electron groups
Seesaw effect
In what would have been a triangular bipyramid structure, the remaining lone pair of electrons exist as a ‘bond’, forming a see-saw shape
See book
T-shape effect
The tendency for atomic bonds to form a T-shape, when left with two electron pairs, full geometric shape would be the ‘triangle bipyramidal geometry’
Square pyramidal
Molecular geometry of what would have been an ocahedron, with the lone pair of electrons at the bottom
Square planar
Molecular geometry of what would have been an ocahedron, with one lone pair of electrons at the bottom and anther lone pair of electrons at the top
Determining the geometry of complex molecules
1) List all of the elements present in the molecule
2) list all of the electron groups for each element type
3) list how many of those electron groups are lone pairs
4) record the best molecular geometry for each element type
Bent
Molecular geometry resulting from a molecule with two lone pairs and two electron groups
Polarity
If the entire molecule where a sphere, would one side of the sphere have more electrons than the other? If so, that side of the molecule has an overall negative charge while the opposite side has a positive charge and the entire molecule is polar.
Valance bond theory
A chemical bond results when valance electrons from one atom are donated to the unfilled orbitals of another atom
Hybridization
Mathmatical proceedure in which the individual orbitals of an atom come together to form new orbitals for the whole atom (hybrid orbitals)
1) the number of standard atomic orbitals added together always equals the number of hybrid orbitals formed.
2) the partial combonation of standard atomic orbitals added together determines the shapes and energies of the hybrid orbitals formed
3) the particular type of hybridization that occurs is the one that yields the lowest overall energy for the molecule
Hybrid orbitals
New orbitals made from the combination of atomic orbitals when an electron is shared between them
sp^3 hybridization
When a 2s and all three 2p sub-shells come together to form four sp^3 subshells, in tetrahedral geometry with 4 electron groups
sp^2 hybridization
A 2s and 2 of the 2p-orbitals come together in triangular planar geometry to form 3 sp^2 orbitals, leaving one p-subshell with 3 electron groups
Sp hybridization
When a 2s and a 2p forms 2 sp orbitals and leaves 2 p-orbitals in linear geometry, with 2 electron groups
Sp^3d hybridization
When a 3s, all three 3p, and a 3d form 5 new sp^3d hybrid orbitals, leaving four 3d orbitals, in triangular bipyramidal geometryo
Sp^3d^2 hybridization
When a 3s, all three 3p, and two 3d form 6 new sp^3d^2 hybrid orbitals, leaving thee 3d orbitals, in octahedral geometry
Structural isomers
Molecules with the same molecular formula but different structures
Dipole moments
A measure of the separation of positive and negative poles in a molecule
Polyatomic molecules
Molecules composed of three or more elements
Radio activity
The emmision of subatomic particles or high energy electromagnetic radiation by the nuclei of an atom
Radioactive
A term reserved for atoms whose nuclei emit subatomic particles
Phosphorescence
Long-lived emmision of light following the absorption of light
Glow-in-the-dark-effect
Types of radio activity
Alpha decay Beta decay Gamma decay Ray emmition Positron emmition
Nuclide
A isotope of a particular variety in terms of an element
Alpha decay
An unstable nucleus emmits particles composed of two protons and two neutrons (He-4)
Nuclear equation
Way of representing alpha decay
(Parent isotope) → (daughter particle)+’He-4’
Ionizing power
The ability of radiation to ionize other molecules, alpha radiation is highest
Penetrating power
Opposite of ionizing power
Ability for radiation to penetrate matter
Alpha radiation particles (He-4) are too big to penetrate matter
Beta decay
An unstable nucleus that emmits electrons
Beta particle
An electron described as isotope notation
0
-1 e
Beta radiation equation
(Old isotope)→(new isotope)+e
Gamma ray emission
Electromagnetic frequency in the gamma range
Symbolized
0
0 γ
Positron emmition
An unstable nucleus emits a positron
Positron
Anti-particle of the electron
0
+1 e
Electron capture
When an nucleus assimilates an electron from an inner orbital of its electron cloud
Strong force
Natural force that binds the nucleus of an atom together
Nucleons
Protons and neutrons that make up a nucleus
N/Z ratio
Ratio of neutrons to protons
High N/Z ratio
Nucliedes lie above the valley of stability
Low N/Z
Nucliedes lie below the valley of stability
Nucleides
Radioactive isotopes
Valley of Stability
Proper ratio of neutrons to protons for stability, individual for each element
Magic number atoms
Atoms that contain
2, 8, 20, 28, 50, 82 neutrons
126 protons
Are always uniquely stable
Film-badge dosimeters
Photographic film encased in plastic and pinned to clothing
Simple radiation detector
Geiger-muller counter
Particles pass through an argon chamber creating a trail of ionized argon atoms
Scintillation counter
Emmitions pass through a material that emmits light in the presence of radio activity
Rate of decay
Rate=k*N
K is the decay rate
Half-life of decay
T sub-1/2= .693/k
k is the decay rate
Integrated rate law
Ln Nt/N0 = -k*t
Nt is the number of radioactive nuclei at a time
N0 is the initial number of radioactive nuclei
Radiocarbon dating
Devised by Willard Libby in 1949
Used to estimate the ages of fossils and artifacts
Convert halflife to ‘k’-rate
Use integrated rate law
Uranium/Lead Dating
Dating method used for non-living objects or objects older than 50,000 years
Nuclear fision
The splitting of the uranium atom
Chain reaction
Neutrons produced by the fission of one uranium nucleus induces the fission in other uranium nuclei
Critical mass
Enough U-235 to produce a self sustaining reaction
Nuclear reactor
Method of electricity production powered by nuclear fision to produce steam
Converting energy to mass
E=mc^2
Mass defect
The phenomenon in which an isotope-particle has a mass less than the sum of their individual components
Nuclear binding energy
The energy required to break a nucleum into individual nuclei
Binding energy curve
Graphical relationship that describes how mass relates to the binding energy per nucleon
Nuclear Fusion
The combonation of two lighter nuclei to form a heavier nuclei
Transmutation
The transformation of one element into another
Linear Accelerator
A device in which a charged particle is accelerated
Cycletron
A charged particle is accelerated back and forth between two chambers of accelerated voltage
Positron emission tomography
The use of positron emitting nucleides as an imaging technique
Radiotracer
Diagnosis technique in which a radioactive nuclide attatched to a compound or introduced into a mixture in order to track the movement of the diagnosis or mixture within the body.
Radiotherapy
The used of radiation to kill rapidly dividing cells
Other uses for radioactivity
1) Kill micro-organisms
2) used to kill bacteria within food
3) used to control the population of harmful insects
Polyprotic Acid
Acids that contain more than one ionizable proton
Diprotic acid
Acid that has both a strong ionizable proton (H+) and a weak second ionizable proton
Particle
A singular unit of matter, whether it be a molecule, atom, ion, nucleon, or electron
Predicting molecular geometry
See you tube videos
‘Coboltus’ prefix
Co+2
‘Ferrous’ prefix
Fe+2
‘Plumbous’ prefix
Pb 2+
‘Manganous’ prefix
Mn 2+
‘Mercuric’ prefix
Hg 2+
‘Nickelous’ prefix
Ni 2+
‘Stannous’ prefix
Sn 2+
‘Cobaltic’ prefix
Co 3+
‘Ferric’ prefix
Fe 3+
‘Manganic’ prefix
Mn 3+
‘Plumbic’ prefix
Pb 4+
‘Stannic’ prefix
Sn 4+
‘Amonium’ prefix
NH4
‘Hydronium’ prefix
H3O+
‘Mercurous’ prefix
Hg2 2+
‘Acetate’ suffix
C2H3O2-
‘Bicarbonate’ suffix
HCO3-
‘Bisulfate’ suffix
HSO4-
‘Chlorate’ suffix
ClO3-
‘Chlorite’ suffix
ClO2-
‘Dihydrogen phosphate’ suffix
H2PO4-
‘Hydroxide’ suffix
OH-
‘Hypochlorite’ suffix
ClO-
‘Nitrate’ suffix
NO3-
Metalic -suffix
-ide
‘Nitrite’ suffix
NO2-
‘Perchlorate’ suffix
ClO4-
‘Permanganate’ suffix
MnO4-
‘Biphosphate’ suffix
HPO4 2-
‘Carbonate’ suffix
CO3 2-
‘Chromate’ suffix
CrO4 2-
‘Dichromate’ suffix
Cr2O7 2-
‘Peroxide’ suffix
O2 2-
‘Sulfate’ suffix
SO4 2-
‘Sulfite’ suffix
SO3 2-
‘Thiosulfate’ suffix
S2O3 2-
‘Phosphate’ suffix
PO4 3-
‘Phosphite’ suffix
PO3 3-
Binary (H+)(_ -) acid
‘Hydro_ic acid’
Oxy acid nomenclature (the anion ends in oxygen)
H+)(_O
- Ate suffix….’_ic acid’
- ite suffix…‘_ous acid’
Cone-shaped bond
Look it up
Dashed bond
Look it up
Partial Pressure
P.solvent=(n-fraction.solvent)*P.solvent
P.solute=(n-fraction.solute)*P.solute
Friability
The ability of a substance to be easily broken down into smaller pieces
