Lecture Notes

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Chapter 3

• Lecture Notes

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Law of Definite Proportion

• The law of definite proportions states that a
  chemical compound always contains the same
  elements in exactly the same proportions by
  weight or mass.
• The law of definite proportions also states that
  every molecule of a substance is made of the same
  number and types of atoms.

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Law of Conservation of Mass

• The law of conservation of mass states that mass
  cannot be created or destroyed in ordinary
  chemical and physical changes.
• The mass of the reactants is equal to the mass of
  the products.

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Law of Conservation of Mass
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Law of Multiple proportions

• The law of multiple proportions states that when
  two elements combine to form two or more
  compounds, the mass of one element that combines
  with a given mass of the other is in the ratio of
  small whole numbers.

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History of the Atom

• Democritus (Greek Philosopher 470-370 B.C.)
  thought all forms of matter was made up of
  invisible particles called atoms. Democritus
  based is idea about matter on observation.
• John Dalton (1766-1855) revises Democrituss
  theory. Daltons Atomic Theory is based upon
  scientific experimentation. He did not know
  about subatomic particles.

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Daltons Atomic Theory

• All matter is composed of extremely small
  particles called atoms, which cannot be
  subdivided, created, or destroyed.
• Atoms of a given element are identical in their
  physical and chemical properties.
• Atoms of different elements differ in their
  physical and chemical properties.

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Daltons Atomic Theory , continued

• Atoms of different elements combine in simple,
  whole-number ratios to form compounds.
• In chemical reactions, atoms are combined,
  separated, or rearranged but never created,
  destroyed, or changed.
• After years of experimenting, Daltons theory
  has been revised. Today, we know that atoms can
  be divided into smaller particles and we can
  create and destroy atoms (Nuclear Chemistry).
  Also like atoms will combine, such as O2.

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Subatomic Particles

• Experiments by several scientists in the
  mid-1800s led to the first change to Daltons
  atomic theory. Scientists discovered that atoms
  can be broken into pieces after all.
• The smaller parts that make up atoms are called
  subatomic particles.
• The three subatomic particles that are most
  important for chemistry are the electron, the
  proton, and the neutron.

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• Joseph John Thomson (1856-1940) discovers the
  electron in 1897. He also describes the atom as
  a bowl of plum pudding a ball of positive
  material (pudding) containing electrons
  (raisins).

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Subatomic Particles

• Ernest Rutherford (1871-1937) discovers the
  nucleus in 1909 and is also credited with
  discovering the proton. He describes the atom as
  electrons orbiting around the nucleus (a ball of
  positive material) like planets orbiting around
  the sun.
• James Chadwick (1891-1974) discovers the neutron
  in 1932.

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Subatomic Particles

• An electron is a subatomic particle that has a
  negative electric charge.
• Protons are the subatomic particles that have a
  positive charge and that is found in the nucleus
  of an atom.
• The number of protons of the nucleus is the
  atomic number, which determines the identity of
  an element.
• Neutrons are the subatomic particles that have no
  charge and that is found in the nucleus of an
  atom.

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Subatomic Particles
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Atomic Number

• The number of protons that an atom has is known
  as the atoms atomic number.
• Atomic numbers are always whole numbers.
• The atomic number also reveals the number of
  electrons in an atom of an element.
• For atoms to be neutral, the number of negatively
  charged electrons must equal the number of
  positively charged protons.

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Mass Number

• The mass number is the sum of the number of
  protons and neutrons in the nucleus of an atom.
• You can calculate the number of neutrons in an
  atom by subtracting the atomic number (the number
  of protons) from the mass number (the number of
  protons and neutrons).
• mass number atomic number number of neutrons

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Atomic Structures Can Be Represented by Symbols

• The atomic number always appears on the lower
  left side of the symbol.

• Mass numbers are written on the upper left side
  of the symbol.

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Isotopes

• All atoms of an element have the same atomic
  number and the same number of protons. Atoms do
  not necessarily have the same number of neutrons.
• Atoms of the same element that have different
  numbers of neutrons are called isotopes.
• One standard method of identifying isotopes is to
  write the mass number with a hyphen after the
  name of an element.
• helium-3 or helium-4

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Isotopes

• The second method of identifying isotopes shows
  the composition of a nucleus as the isotopes
  nuclear symbol.

• All isotopes of an element have the same atomic
  number. However, their atomic masses are not the
  same because the number of neutrons of the atomic
  nucleus of each isotope varies.

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Atomic Models, continued

• Niels Bohr (1885-1962) discovers in 1913 the
  electrons exist in a specific region around the
  atom that are called energy or quantum levels.
• According to Bohrs model, electrons can be only
  certain distances from the nucleus. Each distance
  corresponds to a certain quantity of energy that
  an electron can have.
• An electron that is as close to the nucleus as it
  can be is in its lowest energy level.
• The farther an electron is from the nucleus, the
  higher the energy level that the electron
  occupies.
• The difference in energy between two energy
  levels is known as a quantum of energy.

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Atomic Models, continued

• Bohrs model of an atom
• electrons travel around the nucleus in specific
  energy levels

• Rutherfords model of an atom
• electrons orbit the nucleus just as planets orbit
  the sun

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Atomic Models, continued

• Louis de Broglie (1892-1987) in 1924 that
  electrons have behavior similar to waves. He
  suggested that as waves, electrons could have
  only certain frequencies, which correspond to
  specific energy level in which the electrons are
  found.
• Today, the modern day theory of the atom states
  that electrons have properties of both particles
  and waves and are located in regions around the
  nucleus called orbitals.

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Modern Atomic Model

• The present-day model of the atom takes into
  account both the particle and wave properties of
  electrons.
• In this model, electrons are located in orbitals,
  regions around a nucleus that correspond to
  specific energy levels.
• Orbitals are regions where electrons are likely
  to be found.
• Orbitals are sometimes called electron clouds
  because they do not have sharp boundaries.
  Because electrons can be in other places, the
  orbital has a fuzzy boundary like a cloud.

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Electrons and Light

• By 1900, scientists knew that light could be
  thought of as moving waves that have given
  frequencies, speeds, and wavelengths.
• In empty space, light waves travel at 2.998 ? 108
  m/s (C).
• C ? (wavelength) X f (frequency)
• The wavelength is the distance between two
  consecutive peaks or troughs of a wave.
• The frequency is the number of waves that pass a
  given point in one second.
• The frequency and wavelength of a wave are
  inversely related

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Wavelength and Frequency
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Electrons and Light, continued

• The electromagnetic spectrum is all of the
  frequencies or wavelengths of electromagnetic
  radiation.
• When a high-voltage current is passed through a
  tube of hydrogen gas at low pressure,
  lavender-colored light is seen. When this light
  passes through a prism, you can see that the
  light is made of only a few colors. This spectrum
  of a few colors is called a line-emission
  spectrum.
• Experiments with other gaseous elements show that
  each element has a line-emission spectrum that is
  made of a different pattern of colors.

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Electrons and Light, continued
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Electrons and Light, continued
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Electrons and Light, continued

• In 1913, Bohr showed that hydrogens
  line-emission spectrum could be explained by
  assuming that the hydrogen atoms electron can be
  in any one of a number of distinct energy levels.
• An electron can move from a low energy level to a
  high energy level by absorbing energy.
• Electrons at a higher energy level are unstable
  and can move to a lower energy level by releasing
  energy. This energy is released as light that has
  a specific wavelength.
• Each different move from a particular energy
  level to a lower energy level will release light
  of a different wavelength.

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Electrons and Light, continued

• An electron in a state of its lowest possible
  energy, is in a ground state.
• The ground state is the lowest energy state of a
  quantized system
• If an electron gains energy, it moves to an
  excited state.
• An excited state is a state in which an atom has
  more energy than it does at its ground state
• An electron in an excited state will release a
  specific quantity of energy as it quickly falls
  back to its ground state.

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Quantum Numbers

• The present-day model of the atom is also known
  as the quantum model.
• According to this model, electrons within an
  energy level are located in orbitals, regions of
  high probability for finding a particular
  electron
• To define the region in which electrons can be
  found, scientists have assigned four quantum
  numbers that specify the properties of the
  electrons.
• A quantum number is a number that specifies the
  properties of electrons.

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Electron Configuration

• The principal quantum number, symbolized by n,
  indicates the main energy level occupied by the
  electron.
• Values of n are positive integers, such as 1, 2,
  3, and 4.
• As n increases, the electrons distance from the
  nucleus and the electrons energy increases.

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Electron Configuration

• An electron in a hydrogen atom can move between
  only certain energy states, shown as n 1 to n
  7.
• In dropping from a higher energy state to a lower
  energy state, an electron emits a characteristic
  wavelength of light.

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Electron Configuration

• In 1925 the German chemist Wolfgang Pauli
  established a rule is known as the Pauli
  exclusion principle.
• The Pauli exclusion principle states that two
  particles of a certain class cannot be in the
  exact same energy state.
• This means that that no two electrons in the same
  atom can have the same four quantum numbers.
• The aufbau principle states that electrons fill
  orbitals that have the lowest energy first.
• Aufbau is the German word for building up.

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Electron Configuration

• The arrangement of electrons in an atom is
  usually shown by writing an electron
  configuration.
• Like all systems in nature, electrons in atoms
  tend to assume arrangements that have the lowest
  possible energies.
• An electron configuration of an atom shows the
  lowest-energy arrangement of the electrons for
  the element.

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Electron Configuration

• Quantum levels (N) can be divided into sublevels
  which are called orbitals.
• Regions where the electron is likely to be found.
• Orbitals have different shapes s (spherical) p,
  d, f (dumbbell) and vary in number per quantum
  level.
• Sublevels N2
• Each orbital can hold 2 electrons
• e- 2 N2

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Shapes of s, p, and d Orbitals
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An Electron Configuration Is a Shorthand Notation

• Based on the quantum model of the atom, the
  arrangement of the electrons around the nucleus
  can be shown by the nucleuss electron
  configuration.
• Example sulfur has sixteen electrons.
• Its electron configuration is written as
  1s22s22p63s23p4.
• Two electrons are in the 1s orbital, two
  electrons are in the 2s orbital, six electrons
  are in the 2p orbitals, two electrons are in the
  3s orbital, and four electrons are in the 3p
  orbitals.

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An Electron Configuration Is a Shorthand Notation

• Each elements configuration builds on the
  previous elements configurations.
• To save space, one can write this configuration
  by using a configuration of a noble gas.
• neon, argon, krypton, and xenon
• The neon atoms configuration is 1s22s22p6, so
  the electron configuration of sulfur is
• Ne 3s23p4