Atomic Theory and Structure Flashcards
Week 3 (94 cards)
1
Q
Greek philosopher that proposed the idea that all matter is made up of atoms
A
Democritus - 5th Century B.C.
2
Q
comes from the Greek word “atomos” which means indivisible.
A
Atom
3
Q
Since the atom is too small to be seen even with the most powerful microscopes
A
Scientists’ reliance on models
4
Q
structure or behavior of atoms cannot be clearly seen
A
Even the world’s best microscope
5
Q
British schoolmaster and chemist formulated a precise definition of the indivisible building blocks of matter called “atomos”
A
John Dalton (18th Century)
6
Q
composed of extremely small particles
A
Elements
7
Q
small particles
A
atoms
8
Q
are identical
A
Atom of elements
9
Q
is different from the atom of all the other elements
A
Atom of one element
10
Q
are composed of atoms of more than one element. The relative number of atoms of each
element in a given compound is always the same.
A
Compounds
11
Q
only involve the rearrangement of atoms.
A
Chemical reactions
12
Q
Atoms are
A
not created or destroyed in chemical reactions.
13
Q
The Classical or Dalton’s Atomic Theory
A
- All matter is made up of tiny, indestructible unit particles called atoms.
- The atoms of a given element are all alike.
- During chemical reactions, atoms may combine or a combination of atoms may break down, but the atoms themselves are unchanged.
- When atoms combine to form molecules, they unite in small whole numbered ratios as 1:1, 1:2, 1:3, 2:
14
Q
when two elements combine to form more than one compound, the mass of one element, which combines with a fixed mass of the other element, will always be ratios of whole numbers
A
Law of Multiple Proportions
15
Q
in a chemical reaction mass is neither created nor destroyed
A
Law of Conservation of Mass
16
Q
British scientist who observed many properties of rays, including the fact that the nature of the rays is the same regardless of the identity of the cathode material. He concluded that cathode rays are not waves but are particles with mass and this led to the discovery of electron
A
Joseph John Thomson ( mid 1800’s)
17
Q
a pioneer in modern physics, discovered the electron in 1895, revolutionizing existing theories of atomic structure.
A
SIR JOSEPH JOHN THOMSON
18
Q
recognized for his investigations into the conduction of electricity through gases and physics
A
Thomas
19
Q
which earned him the 1906 Nobel Prize for
A
Gases (Thomson)
20
Q
and for his work on the mathematics of the
electromagnetic- field theory proposed by James Maxwell.
A
Physics (Thomson)
21
Q
measured mass/charge of e-
A
JJ Thomson (1906 Nobel Prize in Physics)
22
Q
a specialized vacuum tube in which images are produced when an electron beam strikes a phosphorescent surface
A
Cathode Ray Tube
23
Q
positive charge spread over the entire sphere
A
Thomson’s Model of the Atom
24
Q
pictured the atom as consisting of small, negatively charged particles of electricity, or electrons, embedded in a heavier, larger sphere of positive electrical charge, like plums in a pudding
A
Joseph John Thomson (20th Century)
25
first direct and compelling measurement of the electric charge of a single electron
Millikan's Experiment
26
1. atoms positive charge is concentrated in the nucleus
2. proton (p) has opposite (+) charge of electron
3. mass of p is 1840 x mass of e- (1.67 x 10-24 g)
Ernest Rutherford's Experimental Design
27
- experimented with a thin gold sheet, bombarded with alpha particles leading to the
discovery of the nucleus
- a British physicist from New Zealand, devised (1911) a solar-system model of the atom
in which electrons orbit a small, heavy, positively charged central core, or nucleus.
ERNEST RUTHERFORD (1908 Nobel Prize
in Chemistry)
28
described the atom as a tiny, dense, positively charged core called a nucleus, in which nearly all the mass is concentrated.
Rutherford's Model of the Atom
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developed the theory that an electron in an atom could gain or lose energy by absorbing or emitting a quantity of energy in which events, it jumps to a higher or lower orbit. As a result, the atom was visualized as a nucleus surrounded by electrons in different orbits
Niels Bohr (1913)
30
Remained relatively unchanged for over 2,000 years. In the late 18th century new discoveries were made that led to a better understanding of atoms and chemistry.
Many scientists since that time
have contributed new evidence
for the Atomic - Molecular
Theory..
The concept of atoms as proposed by
Democritus
31
- is one of many scientists that have given us a better understanding of Atoms.
- In 1913 the Danish physicist NIELS BOHR revised the planetary model by assuming that electrons can occupy only certain orbits
at specific distances from the nucleus.
- Electrons can jump from alow-energy orbit near the nucleus to orbits of higher energy by absorbing energy (green trails).
- When the electrons return to a lower energy level (purple trails), they release the excess energy in the form of radiation of a characteristic wavelength, such as visible light.
NIELS BOHR
32
32
reveals the configuration of the atom’s 11 electrons. The two inner orbits are completely filled with two and eight electrons, respectively. Only the eleventh electron, which occupies the unfilled outer orbit, takes part in
chemical bonding.
A solar-system model of the sodium atom
33
the electrons are pictured as occupying shells, or layers of space, that are centered on a positively charged nucleus
In a modern representation of the sodium atom
34
Emphasizes proton, neutron and electron distribution; does not accurately depict current
accepted model of atomic structure.
Solar system depiction of atomic
structure.
35
are depicted as clouds of negative charge surrounding the nucleus. The density of the small dots is related to the probability of finding an electron at a particular
location.
Electrons
36
electrons orbit the nucleus in stable orbits. Although not a completely accurate model, it
can be used to explain absorption and emission.
- Electrons move from low energy to higher energy orbits by absorbing energy.
- Electrons move from high energy to lower energy orbits by emitting energy.
- Lower energy orbits are closer to the nucleus due to electrostatics.
Bohr model
37
– noted that the number of positive charges increase from atom to atom by a single electronic unit
Henry Moseley (1914)
38
extended to matter the concept that like light, matter must be both a particle and a wave
Louis de Broglie (1916)
39
– developed an equation that relates the wavelength of an electron to its energy which
describes the probability that an electron will be at a certain point in space.
Erwin Schrödinger (1926)
40
a German physicist who developed the uncertainty principle. This principle states
that it is impossible to determine accurately both the momentum and the position of an electron simultaneously
Werner Heisenberg (1927)
41
It was discovered that if positively charged alpha particles (helium nuclei) were sent into light elements such as beryllium (Be), boron (B), or lithium (Li), an unknown radiant or particle would be accelerated out of the other side of that element.
CHADWICK’s Experiment (1932)
42
- states that all matter is composed of small, fast-moving particles called atoms.
- These atoms can join together to form molecules.
The Atomic - Molecular Theory of Matter
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- is composed of atoms.
- Anything that occupies space and has mass.
Matter
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- have a nucleus which contains protons and neutrons.
Atoms
45
- is surrounded by a cloud of electrons.
- a very small fraction of the volume of an atom.
The nucleus
46
the central portion of the atom. It contains the
protons, p + and neutrons, n 0
Nucleus
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the area around the nucleus where electrons are found. Electrons are arranged within the electron cloud in energy levels.
Electron Cloud or Electron orbit
48
The main energy level (MEL 1, 2, 3, 4, 5, 6,7)
the sub-energy level (SEL s, p, d, f) and the orbitals –shape-square)
Energy levels have layers.
49
Show the mass number and atomic number
Give the symbol of the elemen
Atomic Symbol
50
Counts the number of Protons in the nucleus of an atom. All atoms of an element have the
the same number of protons
= identifies the element
Atomic Number
51
= Counts the number of protons and
neutrons in an atom
= atomic number (Z) + number of neutron
Mass Number
52
composed of atoms
Matter
53
have a nucleus which contains protons and neutrons.
Atom
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cloud that surrounds nucleus
Electron
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a very small fraction of the volume of an
atom.
nucleus
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- An atom is neutral
- The net charge is zero
Number of Electron
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Number of protons is equal to
Number of electrons
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positively charged particle found in the nucleus.
Mass = 1 amu. (a.m.u - Atomic Mass Unit)
Proton
58
Atomic number is equal to
Number of electrons
59
particles with no charge. It is found in the
nucleus.
Mass = 1 amu.
Neutron
60
negative charged particle. It is in the orbit
outside the nucleus.
Mass = 1/1836 amu.
Electron
61
Naturally occurring carbon consists of three
isotopes
12C, 13C, and 14C
62
An atom of zinc has a mass number of
65 Zn 30
63
Number of protons in the zinc atom
30
64
Number of neutrons in the zinc atom
35
65
What is the mass number of a zinc isotope with 37 neutrons?
67
66
atoms having the same atomic number but different mass number
Example : 1H1; 2H1; 3H1 and 22Na11 ; 23Na11; 24Na11
Isotopes
67
– atoms having the same mass number but different atomic number
Example: 3H1 and 3He2 ; 14C6 and 14N7
Isobars
68
atoms having the same number of neutrons but different protons
Example : 31P15 and 32Cl16 ; 13C6
and 14C7
Isotones
69
How many protons, neutrons, and electrons are in 14 6C?
6 protons,
8 (14 - 6) neutrons,
6 electrons
70
How many protons, neutrons, and electrons are in 11 6C?
6 protons,
5 (11 - 6) neutrons,
6 electrons
71
are atoms of the same element (X) with different numbers of neutrons in the nucleus
Isotopes
72
The isotopic composition of an element is always expressed on a percentage basis in terms of the relative number of atoms of the various isotopes present.
ISOTOPES AND % ABUNDANCE
73
no. of atoms of a given isotope x 100 total no. of atoms of all isotopes of the element
% Abundance
74
the mass of the isotope relative to the mass.
Isotopic Mass
75
can be determined by a mass spectrometer.
Isotopic abundances
76
- All atoms of an element have the same number of protons but the number of neutrons can vary. Atoms with the same number of protons and differing numbers of neutrons
- Some Isotopes are unstable. The
nucleus of unstable atoms do not
hold together well.
Isotopes
77
Is the process whereby the nucleus of unstable isotopes releases fast moving particles and energy.
Radioactive decay
78
In radioactive elements
particles in the nucleus change form and release energy.
79
when filling orbitals, start with the lowest energy and proceed to the next highest energy level
Aufbau principle
80
- within a subshell, electrons occupy the maximum number of orbitals
possible.
- Electron configurations are using boxes to represent orbitals.
- Each orbital is filled up by a single electron before any pairing can occur.
Hund’s rule
81
Pauli Exclusion Principle
1. No two electrons in an atom may have the same set of four quantum numbers.
2. Two electrons can have the same values of n, l, and ml, but different values of ms.
3. Two electrons maximum per orbital.
4. Two electrons occupying the same orbital are spin paired.
82
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, and 7p.
energy ordering for atomic orbitals
83
-is a space which can be occupied by two electrons
Orbital
84
An orbital’s size and penetration when treated quantitatively
produces the order of filling represented
85
spherical
s orbitals
86
have two lobes separated by a nodal plane.
p orbitals
87
a plane where the probability of finding an electron is zero (here the yz plane).
nodal plane
88
have more complicated shapes due to the presence of two nodal planes
d orbitals
89
If the number of electrons in the s and p sublevels of the outermost energy level is: equal to eight
noble gas
90
If the number of electrons in the s and p sublevels of the outermost energy level is: less than the total number of main energy levels
the element is a metal
91
If the number of electrons in the s and p sublevels of the outermost energy level is: equal or one more than the total number of main energy levels
the element is a semi-metal or metalloid
92
If the number of electrons in the s and p sublevels of the outermost energy level is: greater by two or more than the total number of main energy levels
the element is a non-metal
