8 min read•Last Updated on June 18, 2024
Dalia Savy
Jeremy Kiggundu
Dalia Savy
Jeremy Kiggundu
A cool thing about the periodic table is that it is organized to demonstrate different trends and properties of elements that can be explained by the pattern of electron configurations and the presence of electron-filled orbitals. The periodicity of the periodic table, or its tendency to recur at intervals, can help you estimate the properties of atoms that haven't even been discovered yet.
For the sake of the AP Chemistry exam, rather than only understanding the trends, you should be able to explain why they happen.
In order to fully understand why the trends occur the way they do, it's important to cover the following topics:
As mentioned before, the periodic trends aren't too difficult to grasp since they follow the chronology of the periodic table. It was purposely made to group chemicals of similar properties together.
It is also important to note that the periodic table is divided into 18 columns (called groups) and 7 rows (called periods).
🤨Properties that differ: Going horizontally, each period is organized in order of increasing atomic number. The atomic number, or the number of protons in an atom's nucleus, determines the basic chemical properties of said element. This trend contributes to the differing effective nuclear charge of elements in the same group, which we'll discuss below.
🤝Shared Properties: In each row, the elements have the same number of occupied electron shells.
Let's compare sodium, which is the first element in period 3, with argon, the last element in period 3.
You could see here that both sodium and argon have a total of three occupied electron shells, following the pattern of elements in the same period. However, sodium has 11 protons (represented by its atomic number of 11), and argon has 18 protons (represented by its atomic number of 18).
🤨Properties that differ: One group on the periodic table is organized so that as you move down a group, the number of occupied electron shells increase.
🤝Shared Properties: Every element in one group has the same number of valence electrons in its outermost shell. Because these elements all have the same number of valence electrons, they can bond to other elements in similar ways. In other words, these elements tend to have similar chemical properties.
Some groups on the periodic table have a name, since the elements in a single group have similar properties. For example, the elements in group 18 are called noble gases. All noble gases are generally unreactive due to their high stability.
Let's now compare neon, the second noble gas in group 18, with xenon, the fifth noble gas.
Both neon and xenon have eight valence electrons in their outermost shell, but neon only has two occupied electron shells, while xenon has five occupied electron shells. The fact that they both have a full octet, or eight valence electrons in their outermost shell, makes them both noble gases. Having a full octet makes these elements very stable, and therefore unreactive.
What two subatomic particles make up the nucleus? Protons and neutrons, right? Since neutrons are neutral, protons are the particles that contribute to the positive charge of the nucleus, or the actual nuclear charge (Z).
Now, try to connect this to Coulomb's law which calculates the attraction between two atoms. Each electron orbiting the nucleus experiences both an attraction to the nucleus and a repulsion from the atom's other electrons.
** In order to fully comprehend the forces electrons feel, remember that opposite charges attract. Electrons are negatively charged, and the nucleus is positively charged.**
Electrons that are in the outer shells of an atom may be shielded by the innermost electrons because of the electron-electron repulsion present. In order to accurately represent the nuclear charge of a nucleus, we must account for both the actual nuclear charge and the charge shielded by other electrons (S).
You do not need to know this formula for the AP exam, but it may help you better understand nuclear charge. Effective nuclear charge is equal to the actual nuclear charge (Z) - the charge shielded by other electrons (S).
Let's try to apply the concepts above to the five periodic trends that you should learn and understand for the AP Chemistry exam. The best way to conceptualize this information is to think about it through the concepts we went over above. When in doubt, think about nuclear charge and the periodicity of the periodic table.
The atomic radius of an atom is the distance between an atom's nucleus and its valence electrons.
Going from left to right on the periodic table, the atomic radii get smaller. As you go right, the atomic numbers increase. This means that there is a higher nuclear charge which increases the pull the nucleus has on the electrons. The closer the electrons are to the nucleus, the smaller the distance.
This trend can also be explained by the fact that all elements in a period have the same number of shells. For example, both Li and F have 2 shells, like Na and Ar both have 3 shells.
As you go down a group on the periodic table, the atomic radii increase. This is because the number of occupied shells increases. For example in group 1, Li has 2 occupied shells while Cs has 6 occupied electron shells (similar to the trend explained above with neon and xenon).
The ionic radius is the distance between the nucleus of an ion and the valence electrons of that said ion.
Electronegativity refers to how strongly a nucleus attracts electrons of another atom.
This is because the elements on the right side of the periodic table (such as noble gases in group 18) have more protons in their nuclei, which gives them a greater positive charge. Having a greater nuclear charge makes the nuclei more effective at attracting electrons.
As you go down a group, the atomic size of an atom increases. Therefore, the nucleus of one atom is farther away from the electrons of another atom, and the attraction between the two is weaker.
** Tip - Fluorine is the most electronegative element on the periodic table, with a value of 4.0. Just remember that and try to compare other elements to where fluorine is located on the periodic table.**
Ionization energy is the amount of energy needed to remove the valence electrons of an atom. Since there are often multiple valence electrons, there are multiple ionization energies. The first I.E. is the amount required to remove the most loosely held electron and the second I.E. is the amount required to remove the second most loosely held electron.
Since size decreases across a period, the nucleus and the electrons are more closely attracted to each other. This stronger attraction makes it harder to remove a valence electron. Thus, it takes more energy to do so.
As you go down a group, the amount of occupied electron shells increases. The valence electrons that are farther away are more loosely attracted to the nucleus. Therefore, it takes less energy to remove them.
Therefore, this element has 2 valence electrons!
Electron affinity is the energy change when an electron is added to an atom in the gaseous state.
The more negative the energy, the more energy is released! Electron affinity is typically negative just because an atom releases energy when it gains an electron. However, how negative depends on this trend. You may be able to explain this trend by thinking about electronegativity.
Because of this, you may expect flourine to have the highest magnitude of electron affinity. However, chlorine does! Flourine is too small of an atom and the electrons are so close together that they would repel, which takes energy.
👉 Watch Jacob Jeffries discuss and demonstrate the periodic trends, as well as go over density.
The atomic number is equal to the number protons found in an atom's nucleus. It defines what element an atom is and its place on the periodic table.
Term 1 of 15
The atomic number is equal to the number protons found in an atom's nucleus. It defines what element an atom is and its place on the periodic table.
Term 1 of 15
The atomic number is equal to the number protons found in an atom's nucleus. It defines what element an atom is and its place on the periodic table.
Term 1 of 15
The periodic table is a tabular arrangement of chemical elements, organized based on their atomic number, electron configuration, and recurring chemical properties. Elements are listed in order of increasing atomic number.
Atomic Number: This is the number of protons found in the nucleus of an atom. It defines the identity of an element.
Element Symbol: This is a one or two-letter abbreviation that represents an element on the periodic table. For example, "H" stands for Hydrogen and "O" stands for Oxygen.
Group (Periodic Table): A column in the periodic table that contains elements with similar properties due to having the same number of valence electrons.
In the context of chemistry, groups refer to columns in the periodic table. Elements within these groups share similar chemical behaviors because they have the same number of valence electrons.
Valence Electrons: These are electrons located on an atom's outermost shell and participate in bond formation with other atoms. They're like hands reaching out from one atom to another for bonding!
Alkali Metals: This is Group 1 on the Periodic Table (excluding Hydrogen). Alkali metals are very reactive due to having only one valence electron which they readily give up to achieve a stable electron configuration.
Halogens: This is Group 17 on the Periodic Table. Halogens are highly reactive and often exist in nature as compounds rather than pure elements.
In the periodic table, periods refer to the horizontal rows. Each period corresponds to the number of electron shells an atom of an element in that period has.
Atomic Radius: This refers to the size of an atom. As you move across a period from left to right, atomic radius generally decreases due to increased nuclear charge pulling electrons closer.
Ionization Energy: This is the energy required to remove an electron from a gaseous atom or ion. It generally increases across a period because increasing nuclear charge makes it harder for electrons to escape.
Electronegativity: This measures how strongly atoms attract bonding electrons. It also generally increases across a period because atoms with more protons can attract electrons more strongly.
The atomic number is equal to the number protons found in an atom's nucleus. It defines what element an atom is and its place on the periodic table.
Protons: Positively charged particles found in the nucleus of an atom.
Electrons: Negatively charged particles that orbit around the nucleus of an atom.
Neutrons: Neutral particles (no charge) found in the nucleus of an atom.
Protons are positively charged subatomic particles found within atomic nuclei.
Neutrons: Neutral subatomic particles found within atomic nuclei alongside protons.
Atomic Number: The number of protons in an atomic nucleus which determines its chemical properties and place in the periodic table.
Ion: An atom or molecule with a net electric charge due to loss or gain of one or more electrons.
Effective nuclear charge refers to the net positive charge experienced by an electron in a multi-electron atom. It's not the full nuclear charge because some of this charge is shielded by other electrons in the atom.
Shielding Effect: This refers to the decrease in attraction between an electron and the nucleus in any atom with more than one electron shell.
Atomic Radius: This is half the distance between two nuclei of two identical atoms bonded together. It can be affected by both effective nuclear charge and shielding effect.
Ionization Energy: This is the energy required to remove an electron from a gaseous atom or ion. It's directly related to effective nuclear charge - higher effective nuclear charges generally mean higher ionization energies.
Valence electrons are the outermost electrons in an atom that participate in chemical reactions.
Electron Shell: This is like a layer of an onion, where each shell can hold a certain number of electrons. The outermost shell is where you'll find the valence electrons.
Covalent Bond: This is when two atoms share their valence electrons, kind of like how best friends might share their favorite toys.
Ionization Energy: This is the energy required to remove a valence electron from an atom. It's like how much effort it would take for someone to pry your favorite toy out of your hands!
Noble gases are elements found in Group 18 (VIII A) on the periodic table. They are characterized by full electron shells, making them very stable and unreactive under normal conditions.
Inert Gas: Another name for noble gases due to their low reactivity under normal conditions.
Periodic Table Groups/Families: These are vertical columns on the periodic table where elements share similar properties including reactivity and preferred types of bonds.
Octet Rule: The principle stating that atoms tend to combine so that each has eight electrons (an octet) in its outermost shell, providing stability.
Coulomb's Law describes the force between two charged objects. It states that this force is directly proportional to the product of their charges and inversely proportional to the square of the distance between them.
Electric Charge: This is a fundamental property of matter that can be either positive or negative, with like charges repelling and opposite charges attracting each other.
Inverse Square Law: This law states that a specified physical quantity or intensity is inversely proportional to the square of the distance from the source of that physical quantity.
Force Field: A region around a charged particle within which a force would be exerted on other charged particles or objects.
The atomic radius refers to half of the distance between two nuclei in two adjacent atoms.
Covalent Radius: This is half of the distance between two identical atom’s nuclei when they are joined by a covalent bond.
Van der Waals Radius: This measures half of total distance across an atom when it isn't bonded to another atom but close enough for noticeable interaction.
Ionic Radius: It refers to measure of an atom's ion in a crystal lattice.
Ionic radius refers to measure of an atom's ion in a crystal lattice. Cations (positively charged ions) are typically smaller than neutral atoms, while anions (negatively charged ions) are larger due to changes in electron configuration upon ionization.
Ionization Energy: The energy required to remove an electron from an atom or ion.
Electron Affinity: The amount of energy released when an electron is added to a neutral atom to form a negative ion.
Isoelectronic Ions: These are ions that have the same number of electrons or identical electronic structures.
Electronegativity is a measure of an atom's ability to attract shared electrons in a chemical bond. Elements with high electronegativity tend to pull electrons towards themselves more strongly.
Polarity: A property of molecules resulting from the uneven distribution of charges due to differences in electronegativity; leads to regions of partial positive and negative charge within the molecule.
Ionization Energy: The energy required to remove an electron from an atom or ion; generally increases with increasing electronegativity because these atoms hold their electrons more tightly.
Electron Affinity: The amount of energy released when an electron is added to a neutral atom; elements with high electronegativity typically have high electron affinity as they readily accept additional electrons.
Ionization energy is the amount of energy required to remove an electron from a gaseous atom or ion.
Electron Affinity: This is the amount of energy released when an electron is added to a neutral atom to form a negative ion.
Valence Electrons: These are electrons that reside in the outermost shell surrounding an atomic nucleus and can participate in chemical reactions.
Atomic Radius: This refers to the size of an atom, typically measured by the distance from its nucleus to its outermost electrons.
Quantum tunneling refers to when particles move through a barrier that they theoretically shouldn't be able to pass according to classical physics, due solely to quantum mechanical phenomena.
Wave-Particle Duality: This principle in quantum mechanics suggests that all particles also have properties of waves, which is why they can "tunnel" through barriers.
Quantum Mechanics: The branch of physics dealing with atomic and subatomic systems, providing the framework for understanding phenomena like quantum tunneling.
Heisenberg Uncertainty Principle: A fundamental concept in quantum mechanics stating that it's impossible to simultaneously know both an electron’s position and momentum accurately - this uncertainty allows for phenomena like quantum tunneling.
Electron affinity refers to the amount of energy released when an electron is added to a neutral atom to form a negative ion.
Ionization Energy: The energy required to remove an electron from a gaseous atom or ion.
Ionic Bonding: A type of chemical bond where one atom transfers one or more electrons to another, resulting in attraction between the positively charged ion and negatively charged ion.
Periodic Trends: Patterns that occur within groups and periods on the periodic table including electron affinity, electronegativity, atomic radius, ionization energy, and metallic character.