Chapter 3 Table of Contents Section 1 Substances Are Made of Atoms

Chapter 3 Table of Contents Section 1 Substances Are Made of Atoms
Atoms and Moles Table of Contents Section 1 Substances Are Made of Atoms Section 2 Structure of Atoms Section 3 Electron Configuration Section 4 Counting Atoms

Section 1 Substances Are Made of Atoms
Chapter 3 Objectives State the three laws that support the existence of atoms. List the five principles of John Dalton’s atomic theory.

Chapter 3 Atomic Theory The idea of an atomic theory is more than 2000 years old. Until recently, scientists had never seen evidence of atoms. The law of definite proportions, the law of conservation of mass and the law of multiple proportions support the current atomic theory.

Atomic Theory, continued
Section 1 Substances Are Made of Atoms Chapter 3 Atomic Theory, continued The figure on the right is a more accurate representation of an atom than the figure on the left.

Section 1 Substances Are Made of Atoms Chapter 3 Atomic Theory, continued The Law of Definite Proportions 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.

Section 1 Substances Are Made of Atoms Chapter 3 Atomic Theory, continued The 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.

Law of Conservation of Mass
Section 1 Substances Are Made of Atoms Chapter 3 Law of Conservation of Mass

Law of Conservation of Mass, continued
Section 1 Substances Are Made of Atoms Chapter 3 Law of Conservation of Mass, continued

Section 1 Substances Are Made of Atoms Chapter 3 Atomic Theory, continued The 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.

Law of Multiple Proportions
Section 1 Substances Are Made of Atoms Chapter 3 Law of Multiple Proportions

Dalton’s Atomic Theory
Section 1 Substances Are Made of Atoms Chapter 3 Dalton’s Atomic Theory In 1808, John Dalton developed an atomic theory. Dalton believed that a few kinds of atoms made up all matter. According to Dalton, elements are composed of only one kind of atom and compounds are made from two or more kinds of atoms.

Dalton’s Atomic Theory , continued
Section 1 Substances Are Made of Atoms Chapter 3 Dalton’s Atomic Theory , continued Dalton’s Theory Contains Five Principles 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.

Dalton’s Atomic Theory , continued
Section 1 Substances Are Made of Atoms Chapter 3 Dalton’s Atomic Theory , continued Dalton’s Theory Contains Five Principles, 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. Data gathered since Dalton’s time shows that the first two principles are not true in all cases.

Chapter 3 Objectives State the three laws that support the existence of atoms. List the five principles of John Dalton’s atomic theory.

Chapter 3 Section 2 Structure of Atoms Objectives Describe the evidence for the existence of electrons, protons, neutrons, and describe the properties of these subatomic particles. Discuss atoms of different elements in terms of their numbers of electrons, protons, neutrons, and define the terms atomic number and atomic mass. Define isotope, and determine the number of particles in the nucleus of an isotope.

Chapter 3 Subatomic Particles
Section 2 Structure of Atoms Subatomic Particles Experiments by several scientists in the mid-1800s led to the first change to Dalton’s 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.

Subatomic Particles, continued
Chapter 3 Section 2 Structure of Atoms Subatomic Particles, continued Electrons Were Discovered Using Cathode Rays To study current, J. J. Thomson pumped most of the air out of a glass tube. He applied a voltage to two metal plates, called electrodes, which were placed at either end of the tube. One electrode, called the anode, was attached to the positive terminal of the voltage source, so it had a positive charge. The other electrode, called a cathode, had a negative charge because it was attached to the negative terminal of the voltage source.

Chapter 3 Section 2 Structure of Atoms Subatomic Particles, continued Electrons Were Discovered Using Cathode Rays, continued Thomson observed a glowing beam that came out of the cathode and struck the anode and the nearby glass walls of the tube. He called these rays cathode rays. The glass tube Thomson used is known as a cathode-ray tube (CRT). CRTs are used in television sets, computer monitors, and radar displays.

Chapter 3 Section 2 Structure of Atoms Subatomic Particles, continued An Electron Has a Negative Charge Because the cathode ray came from the negatively charged cathode, Thomson reasoned that the ray was negatively charged. Thomson confirmed this prediction by seeing how electric and magnetic fields affected the cathode ray. Thomson also observed that when a small paddle wheel was placed in the path of the rays, the wheel would turn. This suggested that the cathode rays consisted of tiny particles that were hitting the paddles of the wheel.

Chapter 3 Section 2 Structure of Atoms Subatomic Particles, continued An Electron Has a Negative Charge, continued Thomson’s experiments showed that a cathode ray consists of particles that have mass and a negative charge. These particles are called electrons. An electron is a subatomic particle that has a negative electric charge. Electrons are negatively charged, but atoms have no charge. Atoms contain some positive charges that balance the negative charges of the electrons.

Chapter 3 Section 2 Structure of Atoms Subatomic Particles, continued An Electron Has a Negative Charge, continued Properties of Electrons

Chapter 3 Section 2 Structure of Atoms Subatomic Particles, continued Rutherford Discovered the Nucleus Thomson proposed that the electrons of an atom were embedded in a positively charged ball of matter. His model of an atom was named the plum-pudding model.

Chapter 3 Section 2 Structure of Atoms Subatomic Particles, continued Rutherford Discovers the Nucleus, continued Ernest Rutherford performed the gold foil experiment, which disproved the plum-pudding model of the atom. A beam of small, positively charged particles, called alpha particles, was directed at a thin gold foil. Rutherford’s team measured the angles at which the particles were deflected from their former straight-line paths as they came out of the foil. Rutherford found that most of the alpha particles shot at the foil passed straight through the foil. But very few were deflected, in some cases backward.

Chapter 3 Gold Foil Experiment

Chapter 3 Section 2 Structure of Atoms Subatomic Particles, continued Rutherford Discovers the Nucleus, continued Rutherford reasoned that only a very concentrated positive charge in a tiny space within the gold atom could possibly repel the fast-moving, alpha particles enough to reverse the alpha particles’ direction. Rutherford also hypothesized that the mass of this positive-charge containing region, called the nucleus, must be larger than the mass of the alpha particle. Rutherford argued that the reason most of the alpha particles were undeflected, was that most parts of the atoms in the gold foil were empty space.

Chapter 3 Gold Foil Experiment on the Atomic Level

Chapter 3 Section 2 Structure of Atoms Subatomic Particles, continued Rutherford Discovers the Nucleus, continued The nucleus is the dense, central portion of the atom. The nucleus is made up of protons and neutrons. The nucleus has all of the positive charge, nearly all of the mass, but only a very small fraction of the volume of the atom.

Chapter 3 Section 2 Structure of Atoms Subatomic Particles, continued Proton and Neutrons Compose the Nucleus 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. Because protons and electrons have equal but opposite charges, a neutral atom must contain equal numbers of protons and electrons. Neutrons are the subatomic particles that have no charge and that is found in the nucleus of an atom.

Chapter 3 Section 2 Structure of Atoms Subatomic Particles, continued Proton and Neutrons Compose the Nucleus, continued Properties of a Proton and a Neutron

Chapter 3 Section 2 Structure of Atoms Subatomic Particles, continued Protons and Neutrons Can Form a Stable Nucleus Coulomb’s law states that the closer two charges are, the greater the force between them. The repulsive force between two protons is large when two protons are close together.

Chapter 3 Section 2 Structure of Atoms Subatomic Particles, continued Protons and Neutrons Can Form a Stable Nucleus Protons form stable nuclei despite the repulsive force between them. A strong attractive force between these protons overcomes the repulsive force at small distances. Because neutrons also add attractive forces, some neutrons can help stabilize a nucleus. All atoms that have more than one proton also have neutrons.

Atomic Number and Mass Number
Chapter 3 Section 2 Structure of Atoms Atomic Number and Mass Number Atomic Number Is the Number of Protons of the Nucleus The number of protons that an atom has is known as the atom’s atomic number. The atomic number is the same for all atoms of an element. Because each element has a unique number of protons in its atoms, no two elements have the same atomic number. Example: the atomic number of hydrogen is 1 because the nucleus of each hydrogen atom has one proton. The atomic number of oxygen is 8.

Atomic Number and Mass Number, continued
Chapter 3 Section 2 Structure of Atoms Atomic Number and Mass Number, continued Atomic Number Is the Number of Protons of the Nucleus, continued 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.

Chapter 3 Section 2 Structure of Atoms Atomic Number and Mass Number, continued Atomic Number Is the Number of Protons of the Nucleus, continued The atomic number for oxygen tells you that the oxygen atom has 8 protons and 8 electrons.

Chapter 3 Visual Concepts Atomic Number

mass number – atomic number = number of neutrons
Chapter 3 Section 2 Structure of Atoms Atomic Number and Mass Number, continued Mass Number Is the Number of Particles of the Nucleus, continued 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 Unlike the atomic number, the mass number can vary among atoms of a single element.

Chapter 3 Section 2 Structure of Atoms Atomic Number and Mass Number, continued Mass Number Is the Number of Particles of the Nucleus, continued Example: a particular atom of neon has a mass number of 20. Because the atomic number for an atom of neon is 10, neon has 10 protons. number of protons and neutrons (mass number) = 20  number of protons (atomic number) = 10 number of neutrons = 10

Chapter 3 Section 2 Structure of Atoms Atomic Number and Mass Number, continued Mass Number Is the Number of Particles of the Nucleus, continued The neon atom has 10 protons, 10 electrons, and 10 neutrons. The mass number is 20.

Chapter 3 Visual Concepts Mass Number

Chapter 3 Section 2 Structure of Atoms Atomic Number and Mass Number, continued Atomic Structures Can Be Represented by Symbols Each element has a name, and the same name is given to all atoms of an element. Example: sulfur is composed of sulfur atoms. Each element has a symbol, and the same symbol is used to represent one of the element’s atoms. Atomic number and mass number are sometimes written with an element’s symbol.

Chapter 3 Section 2 Structure of Atoms Atomic Number and Mass Number, continued 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.

Chapter 3 Section 2 Structure of Atoms Atomic Number and Mass Number, continued Atomic Structures Can Be Represented by Symbols Both numbers may be written with the symbol. An element may be represented by more than one notation.

Chapter 3 Section 2 Structure of Atoms Atomic Number and Mass Number, continued Isotopes of an Element Have the Same Atomic Number 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

Chapter 3 Section 2 Structure of Atoms Atomic Number and Mass Number, continued Isotopes of an Element Have the Same Atomic Number, continued The second method of identifying isotopes shows the composition of a nucleus as the isotope’s 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.

Chapter 3 Section 2 Structure of Atoms Atomic Number and Mass Number, continued Isotopes of an Element Have the Same Atomic Number, continued The two stable helium isotopes are helium-3 and helium-4.

Chapter 3 Section 2 Structure of Atoms Atomic Number and Mass Number, continued Isotopes of an Element Have the Same Atomic Number, continued The Stable Isotopes of Lead

Chapter 3 Visual Concepts Isotopes and Nuclides

Chapter 3 Section 3 Electron Configuration Objectives Compare the Rutherford, Bohr, and quantum models of an atom. Explain how the wavelengths of light emitted by an atom provide information about electron energy levels. List the four quantum numbers, and describe their significance. Write the electron configuration of an atom by using the Pauli exclusion principle and the the aufbau principle.

Chapter 3 Atomic Models Rutherford’s Model Proposed Electron Orbits
Section 3 Electron Configuration Atomic Models Rutherford’s Model Proposed Electron Orbits The experiments of Rutherford’s team led to the replacement of the plum pudding model of the atom with a nuclear model of the atom. Rutherford suggested that electrons, like planets orbiting the sun, revolve around the nucleus in circular or elliptical orbits. Rutherford’s model could not explain why electrons did not crash into the nucleus. The Rutherford model of the atom was replaced only two years later by a model developed by Niels Bohr.

Atomic Models, continued
Chapter 3 Section 3 Electron Configuration Atomic Models, continued Bohr’s Model Confines Electrons to Energy Levels According to Bohr’s 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.

Chapter 3 Section 3 Electron Configuration Atomic Models, continued Bohr’s Model Confines Electrons to Energy Levels Rutherford’s model of an atom electrons orbit the nucleus just as planets orbit the sun Bohr’s model of an atom electrons travel around the nucleus in specific energy levels

Chapter 3 Section 3 Electron Configuration Atomic Models, continued Electrons Act Like Both Particles and Waves Thomson’s experiments demonstrated that electrons act like particles that have mass. In 1924, Louis de Broglie pointed out that the behavior of electrons according to Bohr’s model was similar to the behavior of waves. De Broglie suggested that electrons could be considered waves confined to the space around a nucleus. As waves, electrons could have only certain frequencies which correspond to the specific energy levels.

Chapter 3 Section 3 Electron Configuration Atomic Models, continued Electrons Act Like Both Particles and Waves, continued 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.

Chapter 3 Section 3 Electron Configuration Atomic Models, continued Electrons Act Like Both Particles and Waves, continued According to the current model of an atom, electrons are found in orbitals.

Chapter 3 Visual Concepts Orbital

Chapter 3 Electrons and Light
Section 3 Electron Configuration 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  108 m/s. The wavelength is the distance between two consecutive peaks or troughs of a wave. The distance of a wavelength is usually measured in meters.

Electromagnetic Spectrum
Chapter 3 Section 3 Electron Configuration Electromagnetic Spectrum

Electrons and Light, continued
Chapter 3 Section 3 Electron Configuration Electrons and Light, continued The electromagnetic spectrum is all of the frequencies or wavelengths of electromagnetic radiation. The wavelength of light can vary from 105 m to less than 10–13 m. In 1905, Albert Einstein proposed that light also has some properties of particles. His theory would explain a phenomenon known as the photoelectric effect. This effect happens when light strikes a metal and electrons are released.

Chapter 3 Section 3 Electron Configuration Electrons and Light, continued Einstein proposed that light has the properties of both waves and particles. Light can be described as a stream of particles, the energy of which is determined by the light’s frequency. Light is an electromagnetic wave. Red light has a low frequency and a long wavelength. Violet light has a high frequency and a short wavelength. The frequency and wavelength of a wave are inversely related.

Chapter 3 Section 3 Electron Configuration Electrons and Light, continued Light is an electromagnetic Wave, continued The frequency and wavelength of a wave are inversely related. As frequency increases, wavelength decreases.

Wavelength and Frequency
Chapter 3 Section 3 Electron Configuration Wavelength and Frequency

Chapter 3 Section 3 Electron Configuration Electrons and Light, continued Light Emission 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.

Chapter 3 Section 3 Electron Configuration Electrons and Light, continued Light Emission, continued In 1913, Bohr showed that hydrogen’s line-emission spectrum could be explained by assuming that the hydrogen atom’s 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.

Chapter 3 Section 3 Electron Configuration Electrons and Light, continued Light Provides Information About Electrons 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.

Chapter 3 Section 3 Electron Configuration Electrons and Light, continued Light Provides Information About Electrons, continued 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.

Hydrogen’s Line-Emission Spectrum
Chapter 3 Section 3 Electron Configuration Hydrogen’s Line-Emission Spectrum

Chapter 3 Quantum Numbers
Section 3 Electron Configuration 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. The model does not explain how the electrons move about the nucleus to create these regions. 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.

Quantum Numbers, continued
Chapter 3 Section 3 Electron Configuration Quantum Numbers, continued 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 electron’s distance from the nucleus and the electron’s energy increases.

Principal Quantum Number
Chapter 3 Visual Concepts Principal Quantum Number

Chapter 3 Section 3 Electron Configuration Quantum Numbers, continued The main energy levels can be divided into sublevels. These sublevels are represented by the angular momentum quantum number, l. This quantum number indicates the shape or type of orbital that corresponds to a particular sublevel. A letter code is used for this quantum number. l = 0 corresponds to an s orbital l = 1 to a p orbital l = 2 to a d orbital l = 3 to an f orbital

Chapter 3 Section 3 Electron Configuration Quantum Numbers, continued The magnetic quantum number, symbolized by m, is a subset of the l quantum number. It also indicates the numbers and orientations of orbitals around the nucleus. The value of m takes whole-number values, depending on the value of l. The number of orbitals includes one s orbital three p orbitals five d orbitals seven f orbitals

Chapter 3 Section 3 Electron Configuration Quantum Numbers, continued The spin quantum number, indicates the orientation of an electron’s magnetic field relative to an outside magnetic field. The spin quantum number is represented by: A single orbital can hold a maximum of two electrons, which must have opposite spins.

Chapter 3 Section 3 Electron Configuration Quantum Numbers, continued Quantum Numbers of the First 30 Atomic Orbitals

Electron Configurations
Chapter 3 Section 3 Electron Configuration Electron Configurations 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.

Electron Configurations, continued
Chapter 3 Section 3 Electron Configuration Electron Configurations, continued Two electrons can have the same value of n by being in the same main energy level. These two electrons can also have the same value of l by being in orbitals that have the same shape. These two electrons may also have the same value of m by being in the same orbital. But these two electrons cannot have the same spin quantum number. If one electron has the value of 1/2, then the other electron must have the value of –1/2.

Chapter 3 Section 3 Electron Configuration Electron Configurations, continued 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.

Shapes of s, p, and d Orbitals
Chapter 3 Section 3 Electron Configuration Shapes of s, p, and d Orbitals

Chapter 3 Visual Concepts s Orbitals

Chapter 3 Visual Concepts p Orbitals

1s < 2s < 2p < 3s < 3p
Chapter 3 Section 3 Electron Configuration Electron Configurations, continued An Electron Occupies the Lowest Energy Level Available The aufbau principle states that electrons fill orbitals that have the lowest energy first. Aufbau is the German word for “building up.” The smaller the principal quantum number, the lower the energy. Within an energy level, the smaller the l quantum number, the lower the energy. So, the order in which the orbitals are filled matches the order of energies. 1s < 2s < 2p < 3s < 3p

1s < 2s < 2p < 3s < 3p < 4s < 3d
Chapter 3 Section 3 Electron Configuration Electron Configurations, continued An Electron Occupies the Lowest Energy Level Available, continued The energy of the 3d orbitals is slightly higher than the energy of the 4s orbitals. As a result, the order in which the orbitals are filled is as follows: 1s < 2s < 2p < 3s < 3p < 4s < 3d Additional irregularities occur at higher energy levels.

Chapter 3 Section 3 Electron Configuration Electron Configurations, continued An Electron Occupies the Lowest Energy Level Available, continued This diagrams shows how the energy of the orbitals can overlap.

Chapter 3 Section 3 Electron Configuration Electron Configurations, continued 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 nucleus’s 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.

Chapter 3 Section 3 Electron Configuration Electron Configurations, continued An Electron Configuration Is a Shorthand Notation, continued Each element’s 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 atom’s configuration is 1s22s22p6, so the electron configuration of sulfur is [Ne] 3s23p4

Chapter 3 Section 3 Electron Configuration Electron Configurations, continued An Electron Configuration Is a Shorthand Notation, continued Electron orbitals are filled according to Hund’s Rule. Hund’s rule states that orbitals of the same n and l quantum numbers are each occupied by one electron before any pairing occurs. Orbital diagram for sulfur

Chapter 3 Table of Contents Section 1 Substances Are Made of Atoms

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Chapter 3 Table of Contents Section 1 Substances Are Made of Atoms

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