Periodic Trends

Trend for Electronegativity for a group?

It generally decreases down a group. Why? electronegativity decreases because the bonding pair of electrons is increasingly distant from the attraction of the nucleus.

Trend for Electronegativity for a period

It generally increases as you move from left to right across a period. Why? The positively charged protons in the nucleus attract the negatively charged electrons. As the number of protons in the nucleus increases, the electronegativity or attraction will increase.

Metallic character trend for a group

It increase from top to bottom because the outer electrons are further from the nucleus and require less energy for their removal.

Ionization energy trend for a period

It increases from left to right across a period. Why because the electrons are held tighter by the higher effective nuclear charge and the atomic radius decreases, that is, the atom is smaller. The outer electrons are closer to the nucleus and more strongly attracted to the center.

Nonmetallic character trend for a period

It increases from left to right. Why because the tendency to gain electrons increases on moving across a period due to an increase in the nuclear charge and decrease in the atomic size. Hence, the non-metallic character increases across a period.

Nonmetals and their Periodic trends

Nonmetals are further to the right on the periodic table and have high ionization energies and high electron affinities. The atomic radius of a non-metal is generally smaller than the ionic radius of the same element. Nonmetals have more valence electrons and have high electronegativity.

Group 16

Oxygen Group, considered among the main group elements. oxygen(O), sulfur(S), selenium(Se), tellurium(Te), polonium(Po),livermorium(Lv).

Ionic radius trend for a group

The size of an element's this attribute follows a predictable trend on the periodic table. As you move down a column or group, it increases. This is because each row adds a new electron shell.

Examples of endothermic reactions

Thermal decomposition, photosynthesis, denaturing an enzyme.

How to calculate the Effective Nuclear Charge

Zeff( Effective nuclear charge) = Z-S. Z is the number of protons in the nucleus(atomic number). S is the number of electrons between the nucleus and the electron in question (number of non valence electrons)

Group 1

Alkali metals(+1) : H-hydrogen, Li-lithium, Na-sodium, K-potassium, Rb-rubidium, Cs-cesium, Fr-francium.

Group 2

Alkaline Earth Metals (+2) beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra)

What is the Shielding effect

describes the decrease in attraction between an electron and the nucleus in any atom with more than one electron shell. The more electron shells there are, the greater the shielding effect experienced by the outermost electrons.

Non metallic character

relates to the tendency to accept electrons during chemical reactions

Octet rule

that each atom has eight electrons in its valence shell, giving it the same electronic configuration as a noble g

Characteristics of Group 13

As reductants, the group 13 elements are less powerful than the alkali metals and alkaline earth metals. Nevertheless, their compounds with oxygen are thermodynamically stable, and large amounts of energy are needed to isolate even the two most accessible elements—boron and aluminum—from their oxide ores.

Electron configuration

the arrangement of electrons in the orbitals of an atom

Electron affinity

the energy change that occurs when an electron is acquired by a neutral atom.

Ionization energy trend for a group

As you move from top to bottom, the ionization energy increases. Why because electrons are further from the nucleus and thus easier to remove the outermost one and Because outer electrons are further away from the nucleus as you go down a group, they feel less pull from the nucleus, so they are easier to remove.

Endothermic reaction

A reaction that ABSORBS energy in the form of heat

Exothermic reaction

A reaction where energy is transferred to the surroundings.

Group 13

Boron Group The elements are boron (B), aluminum (Al), gallium (Ga), indium (In), thallium (Tl), and nihonium (Nh).

Metalloids in the Periodic Table

Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium, Polonium

Group 14

Carbon Group. any of the six chemical elements that make up this group of the periodic table—namely, carbon (C), silicon (Si), germanium (Ge), tin (Sn), lead (Pb), and flerovium (Fl).

Statement about ionic radius

Cation is smaller because it loses electrons. The electron-electron repulsion is less because the orbital shrinks, so protons and electron attraction is greater. Anion is larger because it gains electrons. The electron-electron repulsion is greater, because the orbital is greater.

Statement about Electronegativity

Elements with high electronegativity have a strong tendency to gain an electron or from electrons from other atoms. Elements with low electronegativity have a strong tendency to lose an electron or electrons from other atoms. The smaller the atom is , the more it will pull electrons towards its nucleus , therefore higher electronegativity.

Statement about Isoelectronic ions

For isoelectronic ions, the greater the positive charge the smaller the radius, the greater the negative charge the larger the radius.

Metals in the Periodic Table

Group 1 except Hydrogen, Group 2,3,4,5,6,7,8,9,10,11,12, Group 13 except Boron, 113, Only Tin and Lead in Group 14, Bismuth in Group 15, Lanthanides and Actinides Series.

The Main Group elements

Groups 1, 2, 13-18

Group 17

Halogen Group. The halogens are located on the left of the noble gases on the periodic table. These five toxic, non-metallic elements of the periodic table and consist of: fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). The most reactive nonmetal group.

Statement about Ionization Energy

High ionization energy- very difficult to remove an electron, non metals and noble gases have high ionization energies. Low ionization energy-very easy to remove an electron and metals have low ionization energies.

Period

Horizontal rows on the periodic table

Nonmetals in the Periodic Table

Hydrogen, Carbon, Nitrogen, Phosphorus, Oxygen, Sulfur , Selenium, Group 17, Group 18.

How to identify isoelectronic species

Isoelectronic species must have the same number of electrons in total. and have the same total number of electrons. To decide if two or more species are isoelectronic: (a) Write the electronic configuration for electrically neutral 2 atoms of the elements.

Nonmetallic character trend for a group

It decreases from top to bottom because due to increase in the atomic size.

Ionic radius trend for a period

It decreases moving from left to right across a row or period. More protons are added, but the outer valence shell remains the same, so the positively charged nucleus draws in the electrons more tightly. But for the nonmetallic elements, the ionic radius increases because there are more electrons than protons.

Atomic Radius trend for a period

It increases from right to left across a period because, Across a period, effective nuclear charge increases as electron shielding remain constant. A higher effective nuclear charge causes greater attractions to the electrons, pulling the electron cloud closer to the nucleus which results in a smaller atomic radius.

Metallic character trend for a period

It increases from right to left. Why because as the effective nuclear charge acting on the valence shell electrons increases across a period, the tendency to lose electrons will decrease. ... Therefore, these electrons can be lost easily.

Atomic Radius trend for a group

It increases from top to bottom because as the atomic number increases down a group, there is again an increase in the positive nuclear charge. However, there is also an increase in the number of occupied principal energy levels. Higher principal energy levels consist of orbitals which are larger in size than the orbitals from lower energy levels. The effect of the greater number of principal energy levels outweighs the increase in nuclear charge and so the atomic radius increases down a group.

Electron Affinity trend for a period

It increases moving from left to right across a period. Why because the electrons added to energy levels become closer to the nucleus, thus a stronger attraction between the nucleus and its electrons.

Electron Affinity trend for a group

It increases moving from top to bottom. Why because the shielding effect increases, thus repulsion occurs between the electrons. This is why the attraction between the electron and the nucleus decreases as one goes down the group in the periodic table.

Characteristics of Group 14

Members of this group conform well to general periodic trends. The atomic radii increase down the group, and ionization energies decrease. Metallic properties increase down the group. Carbon is a non-metal, silicon and germanium are metalloids, and tin and lead are poor metals (they conduct heat and electricity less effectively than other metals such as copper). Despite their adherence to periodic trends, the properties of theses elemets vary greatly. For example, carbon is a non-metal and behaves as such, whereas tin and lead behave entirely as metals.

Metals and their Periodic trends

Metals typically have relatively low ionization energies. The electronegativities of metals are generally low.The electron affinity of metals is lower than that of nonmetals. The atomic radius of a metal is generally larger than the ionic radius of the same element.

Examples of exothermic reactions

Neutralisation, Fuel Combustion, Carbohydrate Oxidation (used in respiration)

Group 15

Nitrogen Group, consisting of nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb), and bismuth (Bi).pnictogens

Group 18

Noble gases, The noble gases are located in the far right of the periodic table and were previously referred to as the "inert gases" due to the fact that their filled valence shells (octets) make them extremely nonreactive. helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and the radioactive radon (Rn). Oganesson (Og) is variously predicted to be a noble gas. Colorless gases, unreactive because of stable electron configuration.

Statement about Atom's Size

Smaller atoms hold their outer electron more tightly, making it harder to remove an electron therefore it takes more energy to remove them. Larger atoms do not hold their outer electrons as tightly, making it more easier to remove an electron, therefore it takes less energy to remove them.

Characteristics of Metals

The bigger the atom, the more reactive. Large atomic radius , smaller ionization energy- easier to take electrons, because they are excellent electron givers. Atomic radius- it increases. Electronegativity- it decreases, Ionization energy-it decreases. Metals are lustrous, malleable, ductile, good conductors of heat and electricity. Other properties include: State: Metals are solids at room temperature with the exception of mercury, which is liquid at room temperature

What are the characteristics of Group 1

The group 1 elements are all soft, reactive metals with low melting points. They react with water to produce an alkaline metal hydroxide solution and hydrogen. Reactivity increases down the group. have relatively low melting points They also have low densities. They are low enough for the first three (lithium, sodium and potassium) to float on water.

Statement about electron affinity

The higher an element´s electron affinity the more stable it becomes when it gains an electron.

Periodic Trends

The location of the elements in the periodic table can be used to predict their relative atomic radii, ionization energy, electron affinity, and electronegativity.

Characteristics of Nonmetals

They are usually typically brittle when solid and usually has poor thermal conductivity and electrical conductivity. Chemically, they tend to have relatively high ionization energy, electron affinity, and electronegativity. They gain or share electrons when they react with other elements and chemical compounds.

Statement about the Main Group Elements

They readily form ions to have an electron configuration like the stable noble gases. Main group elements become isoelectronic with the nearest Noble gases.

Characteristics of Metalliods

They usually look like metals but behave largely like nonmetals. Physically, they are shiny, brittle solids with intermediate to relatively good electrical conductivity and the electronic band structure of a semimetal or semiconductor.

Group 3-12

Transition metals. These elements in groups of the periodic table are called "transition metals". As with all metals, the transition elements are both ductile and malleable, and conduct electricity and heat. The interesting thing about transition metals is that their valence electrons, or the electrons they use to combine with other elements, are present in more than one shell. This is the reason why they often exhibit several common oxidation states. There are three noteworthy elements in the transition metals family. These elements are iron, cobalt, and nickel, and they are the only elements known to produce a magnetic field. Somewhat reactive (variable positive charges)

Characteristics of Group 2

Usually, there is no need to store these elements in oil, unlike the group one elements. For a metal, they tend to have low melting points and low densities. Being a metal, they are obviously good conductors of heat and electricity.

Group

Vertical columns on the periodic table

How to calculate electron configuration

divide the periodic table into sections to represent the atomic orbitals, the regions where electrons are contained. Groups one and two are the s-block, three through 12 represent the d-block, 13 to 18 are the p-block and the two rows at the bottom are the f-block.The symbols used for writing the electron configuration start with the shell number (n) followed by the type of orbital and finally the superscript indicates how many electrons are in the orbital.

What is the Effective Nuclear Charge

e magnitude of the shielding effect is difficult to calculate precisely. As an approximation, we can estimate the effective nuclear charge on each electron.The effective nuclear charge (often symbolized as Zeff or Z*) is the net positive charge experienced by an electron in a multi-electron atom. The term "effective" is used because the shielding effect of negatively charged electrons prevents higher orbital electrons from experiencing the full nuclear charge

Noble gas configuration

electron configuration of noble gases

Valence electrons

electrons in the outermost energy level

Isoelectronic

having the same number of electrons

Periodic Table

is a tabular display of the chemical elements, which are arranged by atomic number, electron configuration, and recurring chemical properties. Elements in the same group (column) have similar chemical properties. Elements close to each other in the same period(row) have similar chemical properties.

Ionic radius

is the radius of a monatomic ion in an ionic crystal structure.

Metallic character

level of reactivity of a metal

Electronegativity

measures the tendency of an atom to attract a shared pair of electrons. An atom's electronegativity is affected by both its atomic number and the distance at which its valence electrons reside from the charged nucleus.

Atomic Radius

one-half the distance between the nuclei of identical atoms that are bonded together

Ionization energy

the energy required to remove an electron from an atom. The first energy refers to the amount of energy required to remove the outermost electron. The second energy refers to the amount of energy required to remove the second outermost electron and so on. The outermost electron is the easiest electron to remove.

Electron Shielding Effect

the reduction of the attractive forces between the positively charged nucleus of an atom and the outermost electrons.