Chapter 2 Flashcards
Atom
The basic structural unit of an element.
The smallest unit of an element that retains the chemical properties of that element.
Consist of 3 primary particles: protons, neutrons, electrons.
Each different element is composed of a different type of atom.
Nucleus
Small, dense, positively charged region in the center of the atom containing protons and neutrons.
Protons
Positively charged particles. (+) In the nucleus. Equal in magnitude of electrons. Defines what the element is
Neutrons
Uncharged particles.
(n)
In the nucleus.
Electrons
Negatively charged particles located outside of the nucleus of an atom.
Move very rapidly in a relatively large volume of space while the nucleus is small and dense.
(-)
Equal in magnitude of protons.
Atomic number (Z)
The number of protons in the atom.
Bottom left hand side.
Mass number (A)
Sum of the number of protons and neutrons.
Bottom top hand side.
Atomic calculation for MASS NUMBER
mass number = number of protons + number of neutrons
Atomic calculation for NUMBER OF NEUTRONS
number of neutrons = mass number - number of protons
OR
number of neutrons = mass number - atomic number
OR
number of neutrons = A -Z
Isotopes
Atoms of the same element having different masses.
They contain the same number of protons, but different numbers of neutrons.
Isotopes of the same element have identical chemical properties.
Some isotopes are radioactive.
Atomic mass
The weighted average of the masses of all the isotopes that make up an element.
Determining atomic mass
Step 1: Convert the percentage to a decimal fraction.
Step 2: Multiply the decimal fraction by the mass of that isotope to determine the contribution of each isotope.
Step 3: Add the mass contributed by each isotope.
Dalton’s Atomic Theory
The first experimentally based theory of atomic structure.
- All matter consists of tiny particles called atoms.
- An atom cannot be created, divided, destroyed or converted to any other type of atom( in a chemical reaction).
- Atoms of a particular element have identical properties (except for isotopes).
- All matter consists of tiny particles called atoms.
- Atoms of different elements have different properties.
- Atoms of different elements combine in simple whole-number ratios to produce compounds (stable combinations of atoms).
- Chemical change involves joining, separating or rearranging atoms.
- 1, 4, 5 and 6 are still regarded as true.
Evidence for electrons
Electrons were the first subatomic particles to be discovered using the cathode ray tube.
Evidence for protons
Protons were discovered by Goldstein.
Same size and charge as electrons, but opposite sign.
1837x heavier than an electron.
Evidence for neutrons
Postulated to exist in the 1920’s, but not demonstrated to exist until 1932.
Almost the same mass as the proton, but zero charge.
Evidence for the nucleus
Initially assumes protons and neutrons were uniformly distributed throughout the atom –> Rutherford’s “Gold Foil Experiment” lead to the understanding of the nucleus - most alpha particles pass through the foil without being deflected. Some particles were deflected, a few even directly back to the source.
Concluded that most of the atom is empty space - the majority of the mass is located in a small, dense region.
What are the two models of the atom?
Thomson
Rutherford
Rutherford’s atom
Tiny, dense, positively charged nucleus of protons surrounded by electrons.
How do we describe the relationship of the electrons to each other and the nucleus?
Use the measurement of particle energy rather than position.
Spectroscopy
Study of information obtained from absorption or emission of light by atoms.
Used to understand the electronic structure.
Electromagnetic radiation
Travels in waves from a source.
Speed of 3.0 x 10^8 m/s.
Wavelength
The distance between identical points on successive waves.
Each wavelength travels at the same velocity but has its own characteristic energy.
Light is propagated (moves) as a collection of ______ waves.
sine
Electromagnetic spectrum
High energy = short wavelength – low energy = long wavelength.
10^-3 – 10^13
Gamma rays > x rays > ultra violet > visible > infarred > microwave > radio waves
Bohr atom
Atoms can absorb and emit energy via promotion of electrons to higher energy levels and relaxation to lower levels.
Energy is emitted upon relaxation is observed as a single wavelength of light.
Spectral lines are a result of electron transitions between allowed levels in the atoms.
Quantization of energy > excited state > relaxation.
Emission spectra
The emission spectrum of hydrogen led to the modern understanding of the electronic structure of the atom.
Electronic transitions
Amount of energy absorbed in jumping from one energy level to a higher energy level is a precise quantity - the energy of that jump is the energy difference between the orbits involved.
Orbit
What Bohr called the fixed energy levels.
Ground state
The lowest possible energy state.
Bohr theory
Electrons are found only in allowed energy levels (orbits).
Atoms absorb energy by excitation of electrons to higher energy levels AND release energy by relaxation of electrons to lower energy levels.
Energy differences may be calculated from the wavelength of light emitted.
Spectral lines are result of electrons transitions between allowed energy levels.
Modern Atomic Theory
Bohr’s model of the atom failed to explain line spectra of atoms with more than one electron.
One major change of Bohr’s model is that electrons do NOT move in orbits.
Atomic orbitals
Regions in space with a high probability of finding an electron.
Electron density
High when electrons move rapidly within the orbital.
Dmitri Medeleev and Lothar Meyer
Independently developed the precursor to our modern periodic table.
They notices that as you list elements in order of atomic mass, there is a distinct regular variation of their properties.
Periodic Law
The physical and chemical properties of the elements are periodic functions of their atomic numbers,
Period
Part of the periodic table.
A horizontal row of elements.
They contain 2, 8, 8, 18, 18, 32, and 32 elements.
Group
Also called families.
Column of elements.
Elements in a particular group or family share many similarities, as in a human family.
Metals
Elements that tend to lose electrons during chemical change, forming positive ions.
Elements found primarily in the left 2/3 of the periodic table.
Properties: high thermal and electrical conductivities, high malleability and ductility, metallic luster, and solid at room temperature.
Nonmetals
Elements that tend to gain electrons during chemical change, forming negative ions.
Elements found in the right 1/3 of the periodic table.
Properties: brittle, powdery solids or gases, opposite of metal properties.
Metalloids
Have properties intermediate between metals and nonmetals.
Atomic number
The number of protons in the nucleus of an atom of an element.
Nuclear charge or positive charge from the nucleus.
The _______________ is the primary factor in understanding how atoms join together to form compounds
electron arrangement
Electron configuration
Describes the arrangement of electrons in atomic orbitals.
Valence electrons
Outermost electrons.
The electrons involved in chemical bonding.
Schroedinger’s equations
Equations that determine the probability of finding an electron in specific region in space, quantum mechanics.
Charge
An electrostatic phenomena - an electrical property.
Opposite charges attract.
Like charges repel.
A neutral atom has the _______ number of protons and electrons.
same
Elements
Fundamental substances that cannot be broken down by chemical means into simpler substances.
92 naturally occurring elements - 118 elements, including human-made.
Represented by a different chemical symbol - one or two letters, derived from the English or Latin name of the element.
Principal energy levels
n = 1, 2, 3, …
The larger the value of n, the higher the energy level and the father away from the nucleus the electrons are.
The number of sublevels in a principal energy level is equal to n… in n = 1, there is one sublevel.
The electron capacity of a principal energy level (or total electrons it can hold) is 2(n)^2… n = 1 can hold 2(1)^2 electrons, which is 2 electrons.
Sublevel
A set of energy-equal orbitals within a principal energy level.
Specify both the _________ and a _________ when describing the location of an electron.
principal energy level
subshell
Subshell
increase in energy: s < p < d < f
Atomic orbital
A specific region of a sublevel containing a maximum of 2 electrons.
Names by their sublevels and principal energy level.
Each type of orbital has a characteristic shape: “s” is spherically symmetrical, while “p” has a shape much like a dumbbell
Aufbau principle
Helps determine the electron configuration.
Electrons fill the lowest-energy orbital that is available first.
Electron configuration rules
Pauli exclusion principle: each orbital can hold up to two electrons with their spins in opposite directions (paired).
Hund’s rule: each orbital in a subshell is half-filled (with one electron) before any becomes orbital becomes completely filled (with two electrons).
Rules for writing electron configurations
Obtain the total number of electrons in the atom from the atomic number.
Electrons in atoms occupy the lowest energy orbitals that are available beginning with 1s.
Fill subshells according to the order.
The s sublevel has one orbital and can hold two electrons.
The p sublevel has three orbitals. The electrons will half-fill before completely filling the orbitals for a maximum of 6 electrons.
The d sublevel has five orbitals. The electrons will half-fill before completely filling the orbitals for a maximum of 10 electrons.
Anion
Negatively charged.
Attracted to anode.
Cation
Positively charged.
Attracted to cathode.
1875 Crookes’s tube
Cathode rays - something leaves the cathode (-) and travels to the anode (+)
J.J. Thompson
1897
The same no matter what gas is in the tube - something that comes from all kinds of atoms. Deflection by an electric field shows that cathode rays consist of negatively charges particles - electrons.
Cathode rays are deflected in magnetic field as well as electric fields.
Measuring the amount of deflection in field of known strength, he calculated the ratio of the mass of the electron to its charge (couldn’t measure either mass or charge separately).
Goldstein’s experiment
Positive particles
1886, observed positive rays using a perforated cathode.
