02.01 Atoms and Elements:

Dalton’s Atomic Model:

The five main ideas in Dalton's hypothesis:

• All matter is composed of extremely small particles called atoms.

  • The word "atom" comes from a Greek word meaning "cannot be cut into smaller pieces." Atoms have their own internal structure that can vary in mass and composition. The types of atoms that make up matter determine its properties.

• Atoms of a given element are identical in size, mass, and other properties.

  • An element is a substance that cannot be broken down into simpler substances. Atoms of a given element are identical.

• Atoms cannot be subdivided, created, or destroyed in chemical reactions.

  • According to the law of conservation of mass, matter cannot be created or destroyed. However, scientists will discover that atoms can be subdivided under unique circumstances.

• Atoms of different elements can combine in simple, whole-number ratios to form chemical compounds.

  • John Dalton's observations eventually led to the law of multiple proportions, which explains that atoms combine to form compounds in simple, whole-number ratios.

  • There's no such thing as a half atom within a compound.

• In chemical reactions, atoms are combined, separated, or rearranged.

  • The bonds between atoms are broken, rearranged, and reformed into new compounds during chemical interactions.

Atomic Discoveries:

How did various scientists add to our knowledge of atomic theory?

• J.J. Thompson:

  • In 1897, after many cathode ray tests, J. J. Thomson discovered smaller, negatively charged particles inside the atom. He called them corpuscles because they were so small, but we now call them electrons.

  • Thomson hypothesized that these corpuscles were scattered within a positively charged atom.

• Ernest Rutherford:

  • In Rutherford's experiments, he made observations about the behavior of positively charged particles as they were radiated onto a piece of gold foil. He noticed some of the particles did not pass straight through the foil as he expected. Some particles scattered, and some bounced back.

  • Rutherford proposed that there must be something big and positively charged in the center of each gold atom to cause the radiation particles to do this. He named this center the nucleus of the atom. He also proposed that the nucleus was composed of two particles, one that was positively charged (proton) and one that was neutral.

• Niels Bohr:

  • After observing the spectrum of colors that were emitted when a gas was excited (electrons gained energy), he believed the different wavelengths of color represented the energy levels an electron could exist within.

  • Bohr proposed that electrons moved in specific orbits around a positive nucleus based on their energy level. He predicted there were many energy levels in which electrons could exist. They are not randomly scattered, as scientists first predicted.

• James Chadwick:

  • Chadwick found the neutron during his experiments in 1932. He and his fellow scientists bombarded beryllium atoms with alpha particles, and an unknown radiation was produced. It was composed of particles with a neutral electrical charge.

  • Chadwick was able to determine that the neutron existed in an atom's nucleus and that its mass was very close to the mass of a proton.

Experimenting with Atoms:

Crooke’s Cathodes: In one of his experimental investigations, Crookes placed an object within the path of the cathode ray. In one of his experimental investigations, Crookes placed an object within the path of the cathode ray. Turn on the cathode ray below, and observe what happens.

Thomson’s Electron: The English physicist Sir J .J. Thomson conducted more experiments using cathode ray tubes.

The beam is affected by the electrically charged rods. Light is not attracted to or repelled by charged objects, so we can conclude that the beam contains small, charged particles that are affected by them. Because the beam was attracted to the positively charged rod and repelled by the negatively charged rod, we can further conclude that these small particles are negatively charged.

The glow is still present even when a different element is used as the negative electrode. We can conclude that the glow is not the property of any one single element, but all elements may share this property. This means atoms of all elements might contain these small, negatively charged particles.

Thomson was able to calculate a mass-to-charge ratio for these small, negatively charged particles. His calculations showed these newly discovered negative particles are much smaller than even the smallest atom.

Rutherford’s Nucleus: Rutherford expected heavy, fast-moving alpha particles to be able to pass straight through a piece of thin gold foil without being affected by the atoms of gold. He assumed that Thomson's plum pudding model was accurate, that matter and mass were evenly distributed throughout the atom. If this model was correct, there should have been nothing to alter the course of the heavy particles flying through the foil.

Modern Atomic Theory:

The atom is the smallest unit of an element that retains the properties of that element. At the center of the atom, there is a very small region called the nucleus. This area makes up most of the mass of the atom.

The area around the nucleus is called the electron cloud. This area makes up the majority of the volume of the atom. The electron cloud is mostly empty space.

An atom is composed of three different subatomic particles: protons, neutrons, and electrons.

The protons are positively charged, and the neutrons have no charge at all. Both particles stay within the nucleus. These particles determine most of the mass of the atom. Because they are tightly packed in the center, the nucleus is also the dense area of the atom.

Electrons are tiny, negatively charged particles. They constantly revolve around the nucleus and cannot enter it. Their mass is so small, they are barely detectable. Each electron moves in its own unique way. Electrons also have the ability to move from one atom to another or throughout the electron cloud of their own atom.

The overall charge of an atom is neutral when it has an equal number of protons (positives) and electrons (negatives). Forces between the nucleus and electrons affect the electrons' position within the atom.

Positive Charge: Protons Win

Negative Charge: Electrons Win

Neutral: Protons = Electrons

02.02 Periodic Table:

Creating the Periodic Table:

The easiest way to organize things is by their similarities.

The periodic table was first described by Dmitri Mendeleev. The foundations were laid down by Democritus, who theorized that substances were made of invisible unbreakable particles called atoms.

The four fundamental elements are fire, earth, air, and water.

Aristotle theorized that the four elements could be mixed to create new materials, which created interest in alchemy, specifically transmuting ordinary metals like lead into gold.

French chemist Antoine Laurent de Lavoisier described an "element" as a chemical substance that could not be broken down into another substance.

John Dalton was able to figure out a way to estimate the weight of each kind of atom.

Scientists kept observing patterns in the list of elements, and eventually in the mid to late 1800s, Dmitri Mendeleev discovered the periodic table.

Periodic Table: Families:

The letter H symbolizes the element Hydrogen.

The atomic number at the top of the box represents the number of protons contained in an atom of that element. Underneath the element's name is the atomic weight.

The horizontal row is called a period. All the elements of a period have the same number of electron shells.

The columns are called groups. In most groups, the elements have the same number of electrons in their outermost electron shell.

Elements can be divided into three groups: metals, nonmetals, and metalloids.

Most elements in the table are metals.

Transition metals can be found in in numerous products such as coins, jewelry, light bulbs, cars, and even some surprising places such as sunscreen and cell phones. Transition metals include groups 3 through 12.

Group 1 elements beginning with lithium and running vertically to francium, are called alkali metals. Group 2 elements beginning with beryllium and running vertically to radium, are called alkaline earth metals.

Elements in groups 1 and 2 are some of the most reactive elements known. Some names are sodium, potassium, and calcium.

Elements in groups 13 through 16 are a mixed bag and are called BCNOs after the first element in each group—boron, carbon, nitrogen, and oxygen—by some researchers.

Average Atomic Mass:

The mass number of the most common isotope of an element can be determined by rounding the atomic mass to the nearest whole number.

Average atomic mass is calculated by taking the weighted average of the masses of all the isotopes of an element, based on their abundance in nature. This is done by multiplying the mass of each isotope by its relative abundance (expressed as a decimal) and then summing these values together.

Reading the Periodic Table:

How to find the number of neutrons in an element:

1. First, round the atomic mass to the nearest whole number. This becomes what we call the mass number.

2. Take the mass number and subtract the atomic number from it.

The result will tell you the number of neutrons in this element. Mass number minus atomic number equals the number of neutrons.

Alkali metals are extremely reactive and likely to pair up with other elements.

Alkaline earth metals can be used in batteries.

Halogens are used in lighting.

Noble gasses are used in neon lights and used in balloons

02.03 Atomic Models:

Finding Electrons:

Quantum Model: a mathematical model based on quantum theory that represents the probable state of electrons within an atom

Niels Bohr proposed that electrons were particles that circled the nucleus in shells that determined their energies and decided that the lowest energy shell of an atom had to have two electrons by looking at the periodic table..

The lowest energy shell is called n equals one.

Very stable noble gasses do not gain or lose electrons readily..

Every atom can be described by four quantum numbers.

n is the principal quantum number and goes up to 7.

The closest shell to the atom nucleus has n= 1.

The angular momentum quantum number is abbreviated with the Greek symbol l.

Heisenberg Uncertainty Principle: you can't know both the position and the velocity or the momentum of a particle at the same time.

Levels of Electrons:

Electrons are located at different places surrounding an atom's nucleus in what are called orbitals.

Electron Arrangements:

An electron configuration is the arrangement of electrons in their energy levels (shells) around the nucleus of the atom and can be used to represent the location of the electrons in all orbitals.

Orbital Notation:

Orbital Notation represents all of the information given by an electron configuration, but it shows the information more visually.

Hund's first rule says that orbitals of equal energy are each occupied by one electron, each spinning in the same direction, before any orbital in that energy level will receive a second electron.

Hund's second rule says that two electrons in the same orbital will spin in opposite directions to reduce the repulsion of the electrons.

02.04 Electromagnetic Radiation:

Behavior of Electrons:

Albert Einstein and Max Planck agreed that energy could have both wavelike and particle-like properties and that that energy was not continuous, but seemed to transfer in packets.

Quanta is defined as a quantity of energy proportional in magnitude to the frequency of the radiation it represents

Feynman's double split experiment demonstrated that electrons seemed to have the properties of both particles and waves.

In his experiment, a stream of electrons was sent through a wall with two small open slits. The electrons hit the backdrop in specific areas like particles, but when the areas were connected, they showed a wavelike interference pattern.

Wavelengths and Frequencies:

Parts of a Wave:

• Crest: the highest point of a wave

• Trough: the lowest point of a wave

• Amplitude: the height or intensity of a wave

Wavelength is the distance between corresponding points on adjacent waves. This distance is consistent throughout the entire sequence within a wave.

Frequency is defined as the number of waves passing by a fixed point in a given amount of time. Frequency is often measured in waves per second.

Wavelength and frequency are inversely related. If the wavelength is increased (becomes longer), the waves become less frequent.

Electromagnetic Radiation:

Electromagnetic radiation (EMR) is a form of energy that moves in a wavelike motion as it travels. It consists of fluctuating electric and magnetic energy fields that oscillate at right angles to each other.

Visible light is one type of electromagnetic radiation.

Invisible types of EMR are radio waves, X-rays, ultraviolet (UV) rays, and gamma rays.

All forms of EMR travel at the speed of light in the vacuum of empty space, where no air or other gas is present. Other types of EMR move through a medium.

Wave Speed:

Electromagnetic energy moves at the speed of light, which is 3.00 × 108 m/s. The speed of light, c, is equal to the product of a wave's frequency, ƒ, and wavelength, λ.

Speed of Light formula: c = λƒ

02.05 Atomic Energy:

Jumping Electrons:

Transitioning electrons release or absorb energy because they are heated or energized.

An emission spectrum is created when this light is analyzed using a spectroscope, a unique set of colored lines called an emission spectrum identifies the element.

Electrons jump to other orbits around the nucleus and fall again due to released energy.

Excited State: when a particle has a higher energy level than its lowest possible ground state

Ground State: the lowest energy state of an atom or other particle

The closer the orbit is to the central nucleus, the lower the electron's potential energy.

The farther the orbit is to the nucleus, the higher the electron's potential energy

As an electron falls, energy equal to the difference between the two orbits is emitted in the form of light

Measuring By Quanta: Planck and Photons:

Photon: is a basic particle representing a quantum of energy, E.

Energy of a Photon formula: E = hƒ

h = Planck's constant of 6.626 × 10−34 J•s

ƒ = the frequency of the energy.

02.06 Periodic Trends Honors:

Nuclear Charges:

Negatively charged electrons repel one another.

Protons (+) and electrons (–) attract each other.

Across the period, the overall nuclear charge increases by one with each additional proton.

Each atom also has an additional electron in its outer energy level.

The number of inner core electrons in the lower energy levels remains the same.

Electrostatic forces: the forces between particles due to their electrical charge

Effective Nuclear Charge: the nuclear charge felt by the outer shell electrons after you have considered the number of shielding electrons that surround the nucleus.

The competing components that affect the effective nuclear charge are:

1. Protons increase down a group, causing a greater attraction between protons and outer shell electrons.

2. The size of atoms and the number of core electrons also increase down a group, thus increasing the shielding effect.

The nuclear charge of outer shell electrons stays relatively the same for each element down a group because the number of additional protons in each element's nucleus is equal to the number of additional shielding inner core electrons..

Attraction of Charges:

Trends in the periodic table all begin with how particles are attracted to one another. The attraction is based on the signs of the charges. The electric charges (e) in an atom are positive or negative. They are depicted as e+ for a proton and e− for an electron.

This is the formula for Coulomb’s law:

• F is the electrical force between the two charges, measured in Newtons

• k is 8.99 × 10⁹(Nm²)/(C² )

• q₁ is the quantity of charge of one particle, measured in coulombs

• q₂ is the quantity of charge of the second particle, measured in coulombs

• d² is the distance between the objects squared, measured in meters

Magnitude of the Charges: As the magnitude of either charge increases, so does the electrostatic force of attraction or repulsion.

Distance Between Charges: As the distance between charges increase, the electrostatic force decreases.

As the distance increases towards infinity, the force decreases to zero.

Periodic Trends: Radius, Ionization, and Electronegativity:

Atomic radius increases moving down the column.

Atomic radius decreases moving across the row from left to right.

Ionization energy is the energy required to remove one electron from an atom, resulting in a positive ion.

Elements with a lower effective nuclear charge require less energy to give up electrons, so their ionization energy is smaller, due to smaller attraction between the outer shell electrons and the positively charged nucleus.

Ionization energy increases from left to right and moving up a group on the periodic table. It generally decreases down a group as the size of atoms increases. This is because it requires less energy to remove an electron when it is in an outer shell much farther from the positively charged nucleus.

Cations are positively charged and result from a loss of electrons.

Anions are negatively charged and result from a gain of electrons.

Metals naturally form cations that are smaller than neutral atoms by losing one or more electrons. When metals lose all the electrons from their outermost energy shell, the remaining electron cloud is smaller.

Nonmetals naturally form anions that are larger than neutral atoms by gaining one or more electrons. When nonmetals gain enough electrons to fill their outermost energy shell, the electrons experience additional repulsive forces between them. (Like charges repel one another.) This causes the electron cloud to spread out and get bigger.

The ionic radii of cations decrease from left to right within metals. The ionic radii of anions decrease from left to right within the nonmetals. Ionic radii increase down a group for metals and nonmetals.

Electronegativity is a measure of the attraction of an atom for the electrons in a chemical bond.

The higher an element's electronegativity, the greater its attraction for electrons in a chemical bond.

Elements with a high effective nuclear charge exert a strong attractive force on their outer shell electrons, so usually they have high electronegativity values.

Fluorine was given the highest electronegativity value, and all other elements were given values by comparing their electronegativity to that of fluorine. Both cesium (Cs) and francium (Fr) have the lowest electronegativity values.

Periodic Trends: Reactivity and Electron Affinity:

The chemical reactivity of metals going down a column increases.

The chemical reactivity of nonmetals going right across a row decreases.

Electron Affinity: the energy involved when a neutral atom gains an electron

When energy is released, the electron affinity value is represented as a negative value. When energy is absorbed, electron affinity is represented with a positive value. An atom that requires energy to force it to accept an electron will not be stable and will lose that electron spontaneously.

Here are the general trends for electron affinity:

• Electron affinities become more negative for each element across a period from group 1 to group 17. This increase in electron affinity across a period relates to the increase in effective nuclear charge. A greater effective nuclear charge results in a stronger attraction for an additional electron.

• The trend for electron affinity is not as obvious moving down groups on the periodic table. In general, we would expect electron affinity values to become less negative (more positive) going down a group, because each electron is being added to a higher energy level farther from the nucleus. Although this is the case for some groups, electron affinity has numerous exceptions.

The alkaline earth metals in group 2 and the nonmetals in group 15 both have electron affinity values close to zero. This is due to the cancellation of opposing forces: electron repulsion within energy levels and the effective nuclear charge.

Notice that the noble gasses in group 18 all have positive electron affinity values. The noble gasses must be forced to gain an electron because they already have a full outer shell. The noble gasses do not form stable negative ions; the additional electron is lost from the atom soon after it is added.