Chapter 1 Flashcards
What is all matter composed of?
atoms which are subatomic particles
protons carry what charge and have what mass?
+1 charge and have a mass of 1 amu/ 1 Da actual charge is 1.6 x 10-19 C
Neutrons have what charge and what mass
neutral and 1 amu
protons and neutrons can be called what?
nucleons
electrons have what charge and mass
-1 charge and negligible mass actual charge is -1.6 x 10-19 C
What is the atomic number (Z)
the number of protons of the atom and gives the element its unique identity
What is the mass number (A)
the total number of neutrons and protons contained in the nucleus
what is an isotope
atoms of the same element that have different numbers of neutrons
What is Deuterium (D or 2H)
An isotope of (1H that has no neutrons) that has one proton and one neutron
what is a cation
ions that carry a net positive charge
what is an anion
ions that carry a net negative charge
polyatomic ions that have positive charges
- represented by roman numerals - the ion with lesser charge uses suffix -ous - the ion with greater charge use suffix -ic
polyatomic ions with oxygen
- use suffix -ite for fewer oxygens - use suffix -ate for more oxygen atoms
polyatomic ions with hydrogen
- one hydrogen is hydrogen (atom name) - two hydrogens is dihydrogen (atom name)
Bohr model of the hydrogen atom:
electrons orbit the nucleus in spherical shells
Electrons closest to the nucleus…
experience the greatest attractive force so it has greater stability and lower energy level
electrons furthest from the nucleus…
have a higher energy level and be less stable
what is the ground state
the lowest possible energy level closest to the nucleus
when is an electron excited
when it absorbs energy equal to the difference between a higher shell’s energy level and the ground states energy level and jumps to the higher-energy shell
when an electron jumps back down to ground state?
energy is emitted in the form of a photon
energy is admitted and absorbed in the form of
electromagnetic radiation
the energy of electromagnetic radiation formula is
E=hf (h = Plank’s constant 6.63 x 10-34 J) (f = frequency of light)
E=hf can also be described as…
E = hc/ wavelength (h = Plank’s constant 6.63 x 10-34 J) (c = speed of light 3.00 x 108 m/s is a vacuum)
equation to calculate the energy level held by an electron at a certain energy level
E = -(R/n2) R = Rydberg constant 2.18 x10-18 J) n = energy level at which the electron is present
Rydberg formula determines the energy level and wavelength of the emitted or absorbed radiation as an electron moves
deltaE=hc/wavelength= R(1/nf2 - 1/ni2) R = Rydberg constant 2.18 x10-18 J) n = energy level at which the electron is present
Moving from a higher to a lower energy level is associated with
emission
moving from a lower level to a higher one
requires energy absorption
Heisenberg uncertainty principle
posits that the more we know about the position of an electron, the less we know about its momentum and vice versa (cannot know both position and momentum of an electron at the same time)
system of quantum numbers
helps describe electrons to an extent
electrons exist in?
orbitals (areas of space where electrons are likely to be located)
each orbital can hold how many electrons
max of 2
principal quantum number (n)
the energy level of the electron (1 or greater) - relates to the row of the table in which the element is found
azimuthal or angular momentum quantum number (l)
describes the sub shell of the principal quantum number where the electron is found and range from 0 to n - 1 l = 0 is the s subshell l = 1 is the p subshell l = 2 is the d subshell l = 3 is the f subshell
magnetic quantum number (ml)
describes the spatial orientation of the orbital in question within its subshell l = 0 holds one orbital l = 1 holds 3 orbitals l = 2 holds 5 orbitals l = 3 holds 7 orbitals
spin quantum number (ms)
describes the spin orientation of the electron which relates to its angular momentum -1/2 or +1/2 2 electrons in same orbital have opposite spin
electron configuration
describes the arrangement of electrons in sub shells around the atomic nucleus ex. A 3d subshell that contains 8 electrons would be denoted as 3d8
Aufbau principle
electrons fill lower-energy orbitals first to create the most stable electron configuration
Hund’s rule
each electron would rather exist in an orbital by itself if possible so we fill each half way first then fill
Element blocks of the periodic table
valence electrons are
the electrons furthest away from the nucleus
periods are also known as
columns are also known as
- rows
- families
division of the periodic table
effective nuclear charge (Zeff)
the attractive force of the nucleus on the atom’s valence electrons
as number of protons increases from left to right, Zeff also increases but moving down the table Zeff decreases
atomic size or radius
inversely related to Zeff so as we go from left to right, atomic size decreases but as we go down, it increases
ionic radius (the radius of a charged speices)
ionic radius of a cation is smaller than the ionic radisu of an anion in comparison to an uncharged form
ionization energy
the energy required to move one valence electron from a neutral atom in a gaseous state
positive value
the first ionization energy is always smaller than the second
directly related to Zeff so it increases from left to right and decreases as we go down columns
electron affinity
denotes the amount of energy released when an electron is added to an atom
atoms that can more readily accept an electron, such as holegns have higher electron affinities.
Directly related to Zeff so increases from L to R and decreases from up to down
exception is noble gases because they posess stable octet
electronegativity
the tendency of an atom to attract electrons that are shared in a chemical bond between 2 atoms
direclty related to Zeff, increases from L to R and decreases from up to down
difference between electron affinity and electronegativity
electronegativity direclty relates to behaviour within a chemical bond
trends summarized
all related to Zeff except atomic size/radius
