periodic table Flashcards
What is the periodic table?
an arrangement of elements in periods and groups and in order of increasing proton number/ atomic number
Describe the change from metallic to
non-metallic character across a period
as we move from left to right across a period in the periodic table, the metallic character of the elements decreases while the non-metallic character increases. This transition is due to the increasing nuclear charge, which results in a stronger attraction between the nucleus and the valence electrons, making it more difficult for the electrons to be removed from the atom and increasing the tendency to form anions.
What is the metallic character related to?
Metallic character is related to the tendency of an element to lose electrons and form cations, while non-metallic character is related to the tendency of an element to gain electrons and form anions.
Describe the relationship between group number
and the charge of the ions formed from elements
in that group
The charge of ions formed from elements in a particular group of the periodic table is related to the number of valence electrons in the outermost shell of the atoms.
In Group 1 of the periodic table, which includes elements like lithium, sodium, and potassium, the atoms have one valence electron. When they react with other elements, they tend to lose this one valence electron to form cations with a charge of +1.
In Group 2 of the periodic table, which includes elements like beryllium, magnesium, and calcium, the atoms have two valence electrons. When they react with other elements, they tend to lose these two valence electrons to form cations with a charge of +2.
In Group 13, which includes elements like boron, aluminum, and gallium, the atoms have three valence electrons. When they react with other elements, they can lose all three valence electrons to form cations with a charge of +3.
In Group 16, which includes elements like oxygen, sulfur, and selenium, the atoms have six valence electrons. When they react with other elements, they tend to gain two electrons to form anions with a charge of -2.
In Group 17, which includes elements like fluorine, chlorine, and bromine, the atoms have seven valence electrons. When they react with other elements, they tend to gain one electron to form anions with a charge of -1.
In general, as we move from left to right across a period of the periodic table, the elements have an increasing number of valence electrons, which tends to make them more likely to gain electrons and form anions. As we move down a group, the number of valence electrons tends to remain the same, but the size of the atoms increases, making them more likely to lose electrons and form cations.
Explain similarities in the chemical properties of
elements in the same group of the Periodic Table
in terms of their electronic configuration
The chemical properties of elements in the same group of the periodic table are similar because they have the same number of valence electrons in their outermost shell. The valence electrons are the electrons that are involved in chemical bonding, and they determine the way in which an atom interacts with other atoms to form compounds.
For example, the elements in Group 1 of the periodic table all have one valence electron in their outermost shell. This means that they all have similar chemical properties when they react with other elements. They tend to lose this valence electron to form ions with a charge of +1, and they are very reactive metals that readily form ionic compounds with non-metals such as halogens.
Similarly, the elements in Group 17 of the periodic table all have seven valence electrons in their outermost shell. This means that they all have a strong tendency to gain one electron to form anions with a charge of -1. They are all highly reactive non-metals that readily form ionic compounds with metals such as Group 1 metals.
The elements in other groups of the periodic table also have similar chemical properties based on the number of valence electrons in their outermost shell. For example, the elements in Group 2 all have two valence electrons and tend to lose these electrons to form ions with a charge of +2, while the elements in Group 16 all have six valence electrons and tend to gain two electrons to form ions with a charge of -2.
In general, the number of valence electrons in an atom determines its chemical properties, and elements with the same number of valence electrons in their outermost shell tend to have similar chemical properties.
Explain how the position of an element in
the Periodic Table can be used to predict its
properties
Atomic structure: The position of an element in the periodic table provides information about its atomic structure. The atomic number of the element is the number of protons in the nucleus of the atom. The number of protons determines the element’s identity and its position in the periodic table. The number of electrons in the outermost shell (valence electrons) determines the element’s chemical properties.
Periods: The elements in the same period have the same number of electron shells. The properties of the elements change across a period from metallic to non-metallic character. Elements in the same period have similar atomic radii and ionization energies.
Groups: The elements in the same group have the same number of valence electrons. Elements in the same group have similar chemical properties, such as reactivity, ionization energy, and electronegativity. The elements in the same group tend to have the same valence electron configuration, and this determines their chemical behavior.
Metals and Non-Metals: Elements in the left-hand side of the periodic table are typically metals, and elements on the right-hand side are typically non-metals. This trend holds true in general, with some exceptions in the transition metals.
Chemical Reactivity: The position of an element in the periodic table can predict its chemical reactivity. The reactivity of elements increases as you move down the periodic table within a group, and decreases as you move across a period from left to right. This is due to changes in the atomic radius, ionization energy, and electronegativity.
In summary, the position of an element in the periodic table can be used to predict its atomic structure, chemical properties, and physical properties. This information can be used to understand how an element will react with other elements and to predict the properties of new compounds.
Describe the Group I alkali metals, lithium,
sodium and potassium
relatively soft metals
with general trends down the group
the Group I alkali metals, lithium, sodium and potassium, exhibit general trends down the group, such as decreasing melting point, increasing density and increasing reactivity. These properties can be explained by the increasing atomic size and the corresponding changes in the electronic configuration of the elements in the group.
What is the general trend down the group of Group 1 alkali metals
(a) Decreasing melting point:
As we move down the group, the melting points of the alkali metals decrease. This is because the metallic bonds between the atoms become weaker, due to the increasing distance between the positively charged atomic nuclei and the negatively charged delocalized electrons. This means that less energy is required to break these bonds and the melting points decrease.
(b) Increasing density:
As we move down the group, the density of the alkali metals increases. This is because the size of the atoms increases as we move down the group, and this causes an increase in the number of atoms per unit volume, resulting in an increase in the density.
(c) Increasing reactivity:
As we move down the group, the alkali metals become more reactive. This is because the outermost electron of each atom is further away from the nucleus, which makes it easier to remove that electron, resulting in a lower ionization energy. Therefore, the alkali metals are more likely to lose this outermost electron and form positively charged ions, which makes them highly reactive.
Predict the properties of other elements in
Group I, given information about the elements
Based on the properties of these elements, we can predict the properties of other elements in the group:
Increasing atomic size: As we move down the group, the atomic size of the elements increases. Therefore, we can predict that the atomic size of the other elements in Group I will also increase down the group.
Decreasing ionization energy: The ionization energy decreases down the group as the outermost electron becomes further from the positively charged nucleus. Therefore, we can predict that the ionization energy of the other elements in Group I will also decrease down the group.
Increasing reactivity: The alkali metals are highly reactive due to their low ionization energy and tendency to lose their outermost electron to form a +1 ion. Therefore, we can predict that the other elements in Group I will also be highly reactive.
Increasing melting and boiling points: The melting and boiling points of the elements in Group I increase down the group due to the increasing strength of metallic bonds. Therefore, we can predict that the melting and boiling points of the other elements in Group I will also increase down the group.
Increasing density: The density of the elements in Group I increases down the group due to the increasing atomic mass and number of atoms per unit volume. Therefore, we can predict that the density of the other elements in Group I will also increase down the group.
In summary, based on the properties of the known elements in Group I, we can predict that other elements in the group will also exhibit similar trends, such as increasing atomic size, decreasing ionization energy, increasing reactivity, increasing melting and boiling points, and increasing density down the group.
Describe the Group VII halogens, chlorine,
bromine and iodine
diatomic non-metals with general trends down the group
Describe the Group VII halogens trends going down the group
(a) Increasing density: As we move down the group, the atomic size of the halogens increases due to the addition of more electron shells. This causes an increase in forces between the atoms, leading to an increase in the density of the halogens. Therefore, the density of chlorine, bromine, and iodine increases down the group.
(b) Decreasing reactivity: As we move down the group, the reactivity of the halogens decreases. This is because the outermost electron is increasingly farther away from the positively charged nucleus, which reduces the attraction between the electron and the nucleus. This leads to a decrease in the ability of the halogens to attract electrons from other atoms or molecules. Therefore, the reactivity of chlorine, bromine, and iodine decreases down the group.
Chlorine is the most reactive halogen, while iodine is the least reactive. Chlorine readily reacts with most metals to form ionic halides, while bromine reacts with fewer metals and iodine reacts with even fewer metals. Iodine is relatively unreactive and is typically only involved in reactions with more reactive elements or compounds.
In summary, the halogens chlorine, bromine, and iodine are diatomic non-metals with general trends down the group of increasing density and decreasing reactivity.
State the appearance of chlorine at r.t.p
a pale yellow-green gas
State the appearance of bromine at r.t.p
a red-brown liquid
State the appearance of iodine at r.t.p
a grey-black solid
Describe and explain the displacement reactions
of halogens with other halide ions
The displacement reactions of halogens with other halide ions occur because halogens are more reactive than other halide ions. In a displacement reaction, a more reactive halogen will displace a less reactive halogen from its compound. For example, when chlorine gas is bubbled through a solution of potassium bromide, the chlorine will displace the bromine from the potassium bromide, and potassium chloride and bromine will be formed.
The reactivity of halogens decreases down the group, so a halogen higher up the group can displace a halogen lower down the group from a compound. For example, chlorine can displace both bromine and iodine from their compounds.
The displacement reactions of halogens with other halide ions can be explained in terms of the relative strength of the covalent bond between the halogen and the halide ion. A more reactive halogen has a stronger bond with the halide ion and can more easily displace a less reactive halogen with a weaker bond.
These displacement reactions are important in the production of certain chemicals, such as chlorine and iodine, and can also be used to test for the presence of halide ions in a sample. For example, adding chlorine water to a solution of an unknown halide salt will produce a precipitate of the halide ion that can be identified by its colour. Chlorine water will produce a white precipitate of silver chloride with a solution containing silver ions, a cream precipitate of silver bromide with a solution containing bromide ions, and a yellow precipitate of silver iodide with a solution containing iodide ions.
