General Chemistry Intro Flashcards
Atoms
All mass consists of tiny particles called atoms
Composition of an atom
Nucleus surrounded by electrons
Necleons: protons and neutrons that make up the nucleus, approximately equal in size and mass
Strong nuclear force
Protons and neutrons are held together to form the nucleus by this force
Stability can be measured by binding energy: energy that would be required to break the nucleus into individual protons and neutrons
Angstrom
One angstrom is 10^-10 m
Radius of nucleus on order of 10^-14 Angstroms
What distance are electrons from the nucleus?
1 to 3 Angstroms
Mass of subatomic particles
Electron: 5.5 x 10^-4 amu
Proton: 1.0073 amu
Neutron: 1.0087 amu
Why is matter mostly empty space?
Matter is composed of atoms
Atom is composed of mostly empty space
Electron Charge
Symbol, e
Equal to 1.6 x 10^-19 Coulombs or C (SI unit for charge)
Charge of an electron or proton
Element
Building blocks of compounds and cannot be decomposed into simpler substances by chemical means
Atomic number, Z
Number of protons, provides identity of element
Each element has a unique number of protons
Mass Number, A
Number of protons plus neutrons
Varies depending on the number of neutrons
Mass number of element is approximately equal to atomic weight or molar mass of element
Isotope
Two or more atoms of the same element that contain different numbers of neutrons
Nucleus of specific isotope is called nuclide
Isotopes have similar chemical properties
Isotopes of Hydrogen
{1}^H, {2}^H, {3}^H
Protium, Deuterium, and Tritium
99.98% of naturally occurring hydrogen is protium
Isotopes of Carbon
{12}^C, {13}^C, {14}^C
6 neutrons, 7 neutrons, and 8 neutrons, respectively
Ion
Number of electrons and protons are not equal in atom
Atom carries a charge and is not electrically neutral
Cations: positively charged, have fewer electrons than protons
Anions: negatively charged, have more electrons than protons
Salt: neutral compound composed of positive and negative ion together
What does changing the number of neutrons, electrons, and protons do to an element?
Creates isotope, creates ion, and changes to another element, respectively
What happens to the atomic radius when a neutral atom loses an electron to become a cation? gains an electron to become an anion?
Atomic radius never changes, ionic radius does
Ionic radius gets smaller because positive charge of the nucleus exerts a greater attractive force on each valence electron, pulling them closer the nucleus
Loss of electron reduces repulsive forces between electrons further contributing to decrease in size
Loss of electron also increases Z_eff for each electron
Ionic radius gets larger
Periodic Table
Table that lists elements from left to right in order of their atomic numbers
Each horizontal row is called a period
Each vertical column is called a group or family
Elements in the same family share similar chemical and physical properties
What are the two methods commonly used to number groups in the periodic table?
Number 1 through 18 left to right is newer
Separate groups into A and B and then number with Roman numerals is older
What are three common groups for elements in the periodic table?
Nonmetals on right, metals in middle and left, metalloids along diagonal separating the metals and nonmetals
Metals
Large atoms that tend to lose electrons to form positive ions and oxidation states
Atoms in a sea of electrons, fluid-like nature of valence electrons
Metallic character: ductility (easily stretched), malleability (easily hammered into thin strips), thermal and electrical conductivity, and luster
Metal atoms easily slide past each other
Electrons move easily from one metal atom to the next, transferring energy or charge (heat or electricity)
All metals except mercury exist as solids at room temperature
Typically lose electrons to become cations (form ionic bonds)
Groups of periodic table that you should know
Group (1): alkali metals, IA Group (2): alkaline earth metals, IIA Group (16): oxygen group, VIA Group (17): Halogens, VIIA Group (18): Noble gases, VIIIA
What are the names of the periods of the bottom two rows of the periodic table?
Lanthanides on top
Actinides on bottom
Nonmetals
Diverse appearances and chemical behaviors
Molecular substances generally made with nonmetals, because tend to covalently bond
Lower melting points than metals
Tend to form anions, reacting with cations to form ionic compounds
Metalloids
Some metallic and some non-metallic characteristics
Representative elements
Section A groups: 1, 2, 13, 14, 15, 16 17, 18
Main-group elements
Make ions by forming closest noble gas electron configuration
Metals tend to form cations and nonmetals tend to form anions to form noble gas config.
Transition Metals
Section B groups (3, 4, 5, 6, 7, 8, 9, 10, 11, 12)
When transition metals form ions, lose electrons from highest s-subshell and then from d-subshell
Common ions formed by transition metals
Group 11 makes +1 ions: Cu+, Ag+, Au+, but also Cu2+, Au3+
Group 6: Cr3+, Group 7: Mn2+, Group 8: Fe2+, Fe3+, Group 9: Co2+, Group 10: Ni2+, Pt2+, Group 12: Zn2+, Cd2+, Hg_2 2+, Hg2+
Group 13: Al3+
Group 14: Sn2+, Pb2+
Group 15: Bi3+
What types of orbitals will an ion have, if possible?
Half-filled or completely filled orbitals
Group 1: half-filled s orbital
Group 2: completely filled s orbital
Group 7: (VIIB) half-filled d orbital
Group 12: (IIB) completely filled d orbital
Group 15: half-filled p orbital
Group 18: completely filled p orbital
Valence electron
Electrons in the outermost shell
Elements in the same group have same number of valence electrons
- Tend to make same number of bonds and exist as similarly charged ions
- Contribute most to an element’s chemical properties
- Located in outermost shell of an atom
- Usually only e-‘s from s and p are considered valence e-‘s
Group 1
Alkali Metals: soft metallic solids with low densities and low melting points
Form 1+ cations (Na+), highly reactive (reacting with most nonmetals to form ionic compounds and hydrogen to form hydrides)
React exothermically (explosively) with water to produce respective metal hydroxide and hydrogen gas
Exist only in compounds in nature
Hydrogen is nonmetal and not like other elements, forms covalent bonds
Hydrogen
Unique and unlike other elements, does not conform to own family
Nonmetal, can form covalent bonds, or ionic bonds with metal cations
Hydrogen is usually colorless, odorless diatomic gas
Acid-base chemistry and intermolecular forces
Group 2
Alkali earth metals: harder, more dense, and melt at higher temperatures than alkali metals
Form 2+ cations (Mg2+)
Less reactive than alkali metals because highest energy electron completes s orbital
Heavier alkaline earth metals are more reactive than lighter ones
Only exist as compounds in nature
Group 14 elements
Form four covalent bonds with nonmetals, but not all are nonmetals
1 nonmetal, 2 metalloids, and 2 metals
All beyond second period form two additional bonds with Lewis bases using d orbitals
Carbon is only element to form strong pi-bonds making double or triple bonds
Group 15 elements
Can form 3 covalent bonds
All beyond second period can form two additional covalent bonds by using d orbitals, and can further bond with Lewis base to form sixth covalent bond
Nitrogen: fourth covalent bond by donating lone pair of electrons to form a bond, forms strong pi-bonds to make double and triple bonds
Phosphorous: can form weak pi-bonds to make double bonds
Group 16 elements
Chalcogens, or oxygen group
Oxygen second most electronegative element, divalent and can form strong pi-bonds to make double bonds and exists as O2 (dioxygen) or O3 (ozone)
Oxygen reacts with metals to form metal oxides and with alkali metals to form peroxides (Na2O2) and super oxides (KO2)
Sulfur: most common form of pure sulfur is yellow solid S8. Metal sulfides (Na2S) are most common in nature. Can form 2, 3, 4, 5, or 6 bonds due to 3d orbital, can form strong double pi-bonds also
Group 17 elements
Halogens
Radioactively stable elements are Fluorine (F2 gas at room temp), chlorine (Cl2 gas at room temp), bromine (Br2 liquid at room temp), and iodine (I2 solid at room temp)
Highly reactive, like to gain electron to attain a noble gas config
In compounds:
- Oxidation states as high as 7+ (except F) when bonding to high electronegative elements
- Hydrogen halides: gaseous hydrogen halides soluble in water (hydrohalic acids)
- Ionic halides: React with metals to form (NaCl)
- F always has oxidation state of -1 in compounds (can only make one bond)
Group 18
Noble Gases (inert gases)
Nonreactive and very stable
Noble gases are normally found in nature as isolated atoms
Gases at room temperature
Common Diatomic Molecules
Hydrogen, Oxygen, Nitrogen, and Halogens
Safe to assume these are in diatomic form unless otherwise stated
Statement “Nitrogen is nonreactive” refers to N2
Four periodic trends
1: atomic radius- increases going down and to left
2: ionization energy- increases going up and to the right
3: electronegativity- increases going up and to the right
4: electron affinity- increases going up and to the right
Atomic radius
Distance from center of nucleus to outermost electron
Corresponds to size of atom
Radius decreases across period, each subsequent element has additional proton which pulls more strongly on surrounding electrons
Moving down, more shells are added, outer electrons are shielded from attraction of protons, so atomic radius decreases going down group
Electrostatic force
Force between charged objects
Attractive between opposite charges, repulsive between like charges
Coulomb’s law
F = k q1 q2 / r^2
F: electrostatic force, q1 and q2, two charges of particles considered, r: distance between two objects
Negative means an attractive force
When calculating for an electron, should use Z_eff and not Z
Effective nuclear charge
Z_eff
Amount of charge felt by most recently added electron
Perfect Shielding: each electron added to atom would be completed shielded from attractive force of all protons in nucleus except for the last proton added, Z_eff would be 1eV for each electron then
Without shielding, Z_eff would equal Z for each electron
How does Z_eff change in periodic table?
Generally increases going left to right and top to bottom
Although energy level of outermost electrons increases down a group, attractive pull of growing positively charged nucleus outweighs additional shielding effects of higher electron shells
Drops from Neon to Sodium because new shell and only one more proton added to outweigh shielding
How can we understand atomic radius trend using Z_eff?
Effective nuclear charge increases from left to right on periodic table, so each additional electron is pulled more strongly toward nuclease
Result is that atoms tend to get smaller when adding electrons across the periodic table
When moving down a group, each drop represents addition of new electron shell, so atoms tend to increase in size moving down a group even though Z_eff increases
Isoelectric Ions
Ions with the same number of electrons, but different elemental identities
E.g. O2-, F-, neutral Ne, Na+, and Mg2+ all have same number of electrons
Electrons feel different Z_eff with diff. # protons
Largest is O2-, and smallest is Mg2+
Ionization energy
Energy needed to detach an electron from an atom
Generally increases from left to right and bottom to top
First ionization energy: energy necessary to remove an e- from a neutral atom in its gaseous state (largest for noble gases)
Second ionization energy: energy required to remove second e- from same atom to form +2 cation (always greater than first bc Z_eff increases after 1 e- is removed)
Remember that as move to right, Z_eff increases, so pulling e- away is harder, however down a trend, distance is more important
Electronegativity
Tendency of atom to attract electrons shared in a covalent bond
When two atoms have diff. Electronegativities, share e- unequally causing polarity
Relative electronegativity determines direction of polarity w/in bond and w/in molecule
Increases across period and up a group
Pauling Scale: ranges from value of 0.79 for Cesium to 4.0 for fluorine
Which 3 E period trends increase going to the right and up?
Energy of Ionization, Electron Affinity, and Electronegativity
What is the electronegativity of a Noble Gas?
Undefined, they tend not to make bonds
Electronegativity provides useful system for predicting which type of bond will form between two atoms
How can electronegativity be used to predict bond type?
Difference of >= 1.6 on Pauling scale form ionic bonds (metals and non-metals)
Moderate differences between 0.5-1.5 on Pauling scale form polar covalent bonds
Small differences (<= 0.4) form nonpolar covalent bonds
Electron Affinity
Willingness of atom to accept an additional electron
Energy released when electron is added to an isolated atom
Tends to increase from left to right and bottom to top, more exothermic
Sign can be diff.- Some require energy to receive e- and some release energy
Electron affinity for noble gases, however, is endothermic, bc they are stable
Quantum Mechanics
Elementary particles can only gain or lose energy in discrete units
Each energy unit is very small, and is only significant when dealing with elementary particles
Bohr Atom
Represents the atom as a nucleus surrounded by electrons in discrete electron shells
Proposed by Niels Bohr
Orbital structure of H atom: single e- orbits H nucleus in electron shell
Four quantum numbers
N, l, m_l, and m_s
Principal quantum number: n, designates the shell level of e-
Azimuthal quantum number: l, designates e-‘s subshell with a distinct shape
Magnetic quantum number: m_l, designates precise orbital within subshell which holds a max of 2 e-
Electron spin quantum number: m_s, -1/2 or +1/2 denotes the spin of the e- in the same orbital
Pauli Exclusion Principle
No two electrons in same atom can have same four quantum numbers
Principal Quantum Number
Designates shell level of electron
Symbol: n
Value: 1, 2, 3, …
Azimuthal Quantum Number
Designates the electron’s subshell, each with a distinct shape, 2nd number
Subshells include s (spherical), p (dumbbell-shaped), d, or f
Symbol: l
Values: 0 (s), 1 (p), …, n-1
Magnetic Quantum Number
3rd Quantum number
Designates precise orbital within a given subshell, 3D orientation
Each orbital holds two electrons
Symbol: m_l
Values: -l to +l
So s subshell is 0, p subshell is -1, 0, or 1
Electron Spin Quantum Number
Fourth quantum number
Designates the spin value of an electron in the exact orbital and subshell
Symbol: m_s
Values: -1/2 or +1/2
Heisenberg Uncertainty Principle
There is an inherent uncertainty in the product of the position of a particle and its momentum
Uncertainty arises from dual nature (wave-particle) of matter
On order of Planck’s constant (6.6 x 10^-34 Js = h)
\delta x \delta p >= h/2
Aufbau principle
“Building up principle”
States that with each new proton added to create a new element, new electron that is added to maintain neutrality will occupy the lowest energy level available
Lower energy state of system, more stable the system
How can you tell the shell level of the most recently added electron?
Representative elements: shell level is given by period
Transition elements: Shell of most recently added electron lags one behind the period
Lanthanides and Actinides: shell of most recently added e- lags two behind period
Orbital shapes
Subshells of s, p, d, and f are orbital shapes
Not true paths that electrons follow, but rather represent probability functions for position of e-
90% chance of finding e- inside given shape
Why are electrons in higher shells at a higher energy level?
Because force is attractive between the proton and electron, work is required to separate them
Force must be applied over a distance
Work is transfer of energy into or out of system (electron and nucleus)
Energy being added to the system by adding e-, increased electrostatic potential eenergy
Electron Conifguration
Lists the shells and subshells of an element’s e-‘s from lowest to highest energy level
Do not have to be written from lowest to highest energy subshells, but usually are
E.g. Na: 1s^2, 2s^2, 2p^6, 3s^1
Fe: 1s^2, 2s^2, 2p^6, 3s^2, 3p^6, 4s^2, 3d^6
What is different about ions formed by transition metals?
Ions are formed by losing electrons from the subshell with the highest principal quantum number first
Generally this is the s subshell
Few exceptions to configurations: Group 6 and 11 have nearly half-filled d subshells
Borrow one e- from highest s subshell so they und with two half-filled subshell
Most likely appears with Cr and Cu
Cr: [Ar] 4s^1 3d^5, Cu: [Ar] 4s^1 3d^10
Hund’s Rule
Electrons will not fill any orbital in the same subshell until all orbitals in that subshell contain at least one electron and that the unpaired electrons will have parallel spins
Paramagnetic Elements
Elements with unpaired electrons
A subshell is not completely filled (e.g. Li)
Spin of each unpaired electron is parallel to the others
Electrons will align with external magnetic field
Diamagnetic Elements
Elements with no unpaired electrons (e.g. He), so subshells are completely filled
Unresponsive to external magnetic field
Emission line Spectrum
Energy emitted when excited electrons of an element fall to lower energy state
Characteristic of a given element
Absorption Line Spectrum
Radiation absorbed when electrons absorb energy to move to a higher energy state
Characteristic of a particular element
How did Max Planck explain absorption and emission spectra?
Electromagnetic energy is quantized and comes only in discrete units related to wave frequency
Energy can only change in discrete increments given by:
\delta E = h f
H is Planck’s constant
Energy of a single photon is given by same equation (according to Einstein)
Electrons can only jump or fall to specific energy levels in atom
Photoelectric Effect
(Einstein) Light is made up of particles called photons
Shining light on metal emits photoelectrons corresponding to intensity and frequency of light
If frequency of light shining on metal is less than necessary quanta of energy, no electrons will be emitted
Intensity above required frequency increases number of emitted photoelectrons, but not KE
Electrons must be ejected by one-to-one photon to electron collisions rather than by combined energies of two or more photons
Minimum energy required to eject e- is work function, Phi
KE = hf - Phi
Covalent Bonds
Nuclei share a pair of electrons, forms a molecule
Known as intramolecular bonds (bonds within molecules)
Formed only between nonmetal elements, predominant in organic chemistry
Attractive and repulsive forces balance each other
Bond length
Distance between nuclei of two atoms in a bond when they are at their lowest possible energy state
If atoms are separated by an infinite distance, forces between them and so energy of bond goes to zero
Atoms closer than bond length (internuclear distance) have a great spike in energy
Energy necessary for complete separation of bond is given by vertical distance on graph between energy at bond length and zero (internuclear distance v potential energy)
Bond Dissociation Energy
Bond energy
Energy necessary for complete separation of bond
Can be seen by plotting potential energy and internuclear distance for two atoms
Vertical distance on graph between lowest stable energy and x axis
Partial Ionic Character
Difference in electronegativity between two atoms in bond is significant, but not large enough to be ionic (polar covalent)
Ionic Bonds
Electronegativities differ vastly between two atoms in a bond
Electron was actually transferred to other atom and oppositely charged ions are held together by electrostatic forces
Ionic compound, separate distinct unit cannot be separated
Lattice structure within a compound
Intermolecular Forces
Molecules interact with each other through intermolecular forces
Much weaker than intramolecular forces of bonds, similarly influenced by charge and electronegativity
Dipole Moment
Center of positive charge in bond does not coincide with center of negative charge, partial positive / partial negative character
Analogous to center of mass
Vector pointing from center of positive charge to center of negative charge
Measured in units of debye, D, given by equation: /mu = q d
Where q is magnitude of charge, d is distance between centers of charge
Polar Bond
Bond that has a dipole moment is polar, but molecule with polar bonds may not have net dipole moment, and therefore may not be polar (symmetry)
Bond without dipole moment is nonpolar
Intermolecular Attractions
Attractions between separate molecules occur due to dipole moments
Weak electrostatic bonds, about 1% as strong as covalent bonds
Attraction roughly proportional to their dipole moments; stronger the dipole, stronger the attraction
Hydrogen Bond
Strongest type of dipole-dipole interaction
Occurs between hydrogen covalently bonded to fluorine, oxygen, or nitrogen (highly electronegative) and a fluorine, oxygen, or nitrogen from another molecule
Large dipole moments cause positive H to attract to negative electronegative elements
Responsible for high boiling point of water
Still much weaker than covalent bond
Induced Dipole
Dipole moment is momentarily induced in otherwise nonpolar molecule
or bond by polar molecule, ion, or electric field
Partial or full charge of polar molecule or ion attracts or repels electrons of nonpolar molecule, separating centers of positive and negative charge
Weaker than permanent dipoles
Instantaneous Dipole
Can arise spontaneously in an otherwise nonpolar molecule
Occur due to movement of electrons and at any given moment may not be distributed exactly between two bonding atoms
Can create induced dipole in neighboring molecule
Weakest and short-lived
Known as London dispersion forces or Van der Waals’ forces
All molecules exhibit London dispersion forces
Do molecular crystals and ice experience intermolecular forces?
Yes, all dipole and intermolecular forces apply, especially when atoms are in such close contact
Generally insignificant in gases because molecules are spread far apart
Naming of Ionic Compounds
Named after cation or anion
Transition metals with different oxidative states: roman numeral
Copper(I): Cu+1, Copper(II): Cu+2 or cuprous and cupric, respectively
Nonmetal: cation ends in -ium, NH4+ is ammonium
Monatomic anions and simple polyatomic have suffix ‘ide’, Hydride for H-
Polyatomic anions w/ multiple Oxygens have suffix ‘ite’ or ‘ate’ relative # O’s
Nitrite: NO2-, Nitrate: NO3-
Hypo and per used for more O’s, hypo-ite, ite, ate, per-ate
Oxyanions with hydrogen have hydrogen in front
For compounds: Put cation in front of anion and for two atoms begin with element to left down and for second element add suffix ‘ide’, Number prefix used for element with more than one atom, dinitrogen tetroxide (N2O4)
Naming Acids / Bases
Named for anions
If anion ends in ‘ide’, acid name starts with ‘hydro’ and ends in ‘ic’, e.g. hydrosulfuric acid (H2S)
For oxyacid: ending ‘ic’ used for multiple O’s, ‘ous’ for fewer O’s, e.g. sulfuric acid H2SO4 and sulfurous acid H2SO3
What are the seven base SI units you should know?
Mass: Kilogram (Kg) Length: Meter (m) Time: Second (s) Electric current: Ampere (A) Temperature: Kelvin (K) Luminous intensity: Candela (cd) Amount of substance: Mole (mol)
What is the derived units for the Newton?
1 N = 1 kg m / s^2
Compound
Substance made from two or more elements in fixed proportions
Empirical Formula
Smallest ratio of whole numbers that can be used to represent proportions of elements in a compound
Ionic compounds are represented by empirical formula
Molecular Formula
Commonly used for molecular compounds, represents the exact number of elemental atoms in each molecule
Glucose: C6H12O6 is molecular formula, CH2O is empirical formula
Percent Composition by Mass
Multiple an element’s atomic weight by number of atoms it contributes to empirical formula and divide by total weight of all atoms
yields mass fraction of that element in the compound (multiple by 100% to get percentage)
Physical Reactions
Reactions of compounds that maintain their molecular structure (and identity)
Examples are melting, evaporation, dissolution, rotating of polarized light
Chemical Reaction
Compound undergoes a reaction and changes its bonding or structure to form a new compound
Examples are combustion, metathesis, and redox reactions
Combination (Synthesis) Reaction
A + B -> C
E.g. Fe(s) + S(s) -> FeS(s)
Decomposition Reaction
C -> A + B
E.g. 2Ag2O(s) -> 4Ag(s) + O2(g)
Single Displacement Reaction
A + BC -> B + AC
Mg(s) + 2HCl(aq) -> MgCl2(aq) + H2(g)
Double Displacement Reaction
AB + CD -> AD + CB
E.g. HCl(aq) + NaOH(aq) -> NaCl(aq) + H2O(l)
Often occur between ionized salts dissolved in water
Helpful to know the charges of common ions to predict new compound’s molecular formula
Ions will often combine to form a new compound with net charge of 0
Redox Reaction
E.g. 2Au3+ 3Zn -> 2Au + 3Zn2+
Transfer of electrons, oxidation or reduction
Combustion Reaction
E.g. C6H12 + 9O2 -> 6CO2 + 6H2O
Usually involves input heat (burning)
Bronsted-Lowry acid-base
HI + ROH -> I- + ROH2+
Lewis acid-base
Ni2+ + 6NH3 -> Ni(NH3)6 2+
What does the delta symbol above reaction arrows indicate?
Heat was added
Theoretical Yield
Amount of product that should be created when a reaction runs to completion, based on stoichiometry
Amount of product created by real experiment is actual yield
Percent Yield = 100 x (Actual yield / Theoretical yield)
Stoichiometry
Determining quantities of products and reactants in chemical equation
Use grams, amu, moles
Mole
6.011 x 10^23 of something
This is Avogadro’s number
12^C serves as standard for this, Avogadro’s number is equal to the number of Carbon atoms in 12 grams of 12^C
6.022 x 10^23 amu = 1 gram
Moles = grams / atomic or molecular weight
Radioactive decay
Process in which atoms spontaneously break apart
All atoms other than hydrogen are subject to some type of spontaneous decay
Atoms with relatively high decay are said to be radioactive, rate at which decay occurs varies dramatically
Atomic nuclei are held together by strong nuclear force, without which they would repel one another
Half-life
Length of time necessary for one half of a given amount of a substance to decay
Radioactive decay follows first order kinetics
Amount of atoms that remain after decay can be expressed as follows:
At = A0 e^(-kt) OR ln(At/A0) = -kt
At is amount at time t, A0 is original amount, k is rate constant, t is time
Exponential Decay
The relationship governing radioactive decay are described in terms of an exponential relationship
Plotting logarithm of amount of atoms as a function of time would produce a straight line semi-log plot
Three types of radioactive decay on MCAT
Alpha decay, Beta decay, and gamma decay
Positron emission and electron capture are both types of beta decay
Alpha Decay
Loss of an alpha particle
Alpha particle is a helium nucleus (contains 2 protons and 2 neutrons)
E.g. {238}^ {92}_ U -> {4}^ {2}alpha + {234}^ {90} Th
Beta Decay
Breakdown of a neutron into a proton and electron, and expulsion of newly created electron
Since a neutron is destroyed, but a proton is created, mass number stays the same, but atomic number increases by one
E.g. {234}^ {90}Th -> {234}^ {91} Pa + {0}^ {-1}_ e + neutrino
Neutrino is virtually massless particle, typically represented by greek letter nu (v)
Positron Emission
Emission of a positron when a proton becomes a neutron
Type of beta decay
Positron can be thought of as electron with positive charge, where both electrons and positrons are considered beta particles
E.g. {22}^ {11}_ Na -> {0}^ {1}_ e + {22}^ {10}_ Ne + neutrino
Electron Capture
Capture of electron and merging of electron with proton to create a neutron
Proton destroyed and a neutron created
E.g. {201}^ {80}_ Hg + {0}^ {-1}_ e -> {201}^ {79}_ Au + {0}^ {0}_ gamma
Where gamma is gamma ray
Gamma ray
High frequency photon
Has no mass or charge, does not change identity of atom from which it is given off
Gamma decay
AKA gamma ray emission
Often accompanies other types of radioactive decay
Can occur when an electron and positron collide:
{0}^ {-1}_ e + {0}^ {1}_ e -> {0}^ {0}_ gamma + {0}^ {0}_ gamma
Matter-antimatter collision called annihilation where mass is destroyed and converted to energy in the form of gamma rays
