HL CHEM PERIODICTY

electrostatic attraction between the positively charged nucleus and the negatively charged electrons within an atom. The strength of this attraction depends on two factors:

-The distance between the nucleus and the electrons (the atomic radius). -The number of shielding electrons within the atom (shielding electrons are inner electrons that tend to 'shield' the outer electrons from the full attraction of the nucleus). Because of these two factors, valence electrons which are found in the outermost energy levels do not feel the full attraction from the protons in the nucleus. The attraction felt by the valence electrons, known as the effective nuclear charge, is less than the actual nuclear charge of the atom.

period 1

1 energy level = contains the elements hydrogen and helium.

IE 1 trends on for down a group

1) distance of the outermost electron from the nucleus 2) electron shielding as you go down a group atomic radii increases therefore distance increases there shielding increases therefore IE decreases and outer e is easier to move

crystal lattice

A 3-dimensional geometric arrangement of the atoms or molecules or ions composing a crystal

diatomic molecule

A molecule consisting of two atoms

displacement reaction

A reaction in which a more reactive element displaces a less reactive element from an aqueous solution of the latter's ions.

Negatively-charged ions are larger than their parent atoms.

By gaining electrons, they have more electrons than protons, which decreases the attraction between the nucleus and the valence electrons. This weaker attraction causes the ionic radius to increase. These extra electrons also increase the repulsion between electrons, which also contributes to the increase in ionic radii. Across a period the ionic radii of negative ions decrease as the nuclear charge increases. Therefore, going from P3− to Cl−, we see a decrease in the ionic radius.

trends of atomic radii down a group

Each additional period means that the outer electrons occupy a main energy level further from the nucleus, therefore, the atomic radius increases down a group.

The decrease in ionisation energy from beryllium to boron can be explained by examining their electron configurations: Beryllium has the electron configuration 1s2 2s2 Boron has the electron configuration 1s2 2s2 2p1

Electrons in p orbitals are of higher energy and further from the nucleus than electrons in s orbitals, therefore they require less energy to remove. The same explanation can be applied for the drop in ionisation energy from magnesium to aluminium, except that the electron configurations are 1s2 2s2 2p6 3s2 and 1s2 2s2 2p6 3s2 3p1 respectively.

alkali metal and halogen relationship

Fluorine is the most reactive halogen, while iodine (or, strictly speaking, astatine) is the least reactive. The halogens react with elements in group 1 to produce salts (the name halogens meaning 'salt formers'). For example, sodium reacts with chlorine to produce the salt sodium chloride according to the following equation: 2Na (s) + Cl2 (g) → 2NaCl (s)

The electron configuration of magnesium is 1s2 2s2 2p6 3s2. From this, deduce the group, block and period that magnesium occupies in the periodic table.

From the electron configuration, we can deduce that magnesium is in group 2 (as it has two electrons in its valence shell). It is an s-block element (its valence electrons are in the s sub-level) and it is in period 3 of the periodic table (it has three occupied main energy levels).

Going down a group in the periodic table, electronegativity decreases.

Group 1 and group 17 elements, for example, show clear decreasing values in electronegativity. The decrease in electronegativity is due to the increase in atomic radius. As we descend the group there is an increasing distance between the nucleus and shared pairs of electrons, therefore the attraction for these electrons decreases. The increase in nuclear charge down a group is counteracted by the increased shielding caused by the extra occupied main energy levels within the atom.

alkali metals

Group 1, 1 electron in outer level, very reactive, soft, silver, shiny, low density; Lithium, Sodium, Potassium, Rubidium, Cesium, Francium The most reactive group of metals - react strongly with cold water.

noble gases

Group 18 lack of reactivity because of full valence shell A group of very unreactive monatomic gases.

tranistion metals

Groups 3-12 on the periodic table; these elements behave increasingly like nonmetals the farther right on the periodic table they lie. They form positive ions with various charges. Relatively stable metals with only moderate reactivity - useful for construction.

explain the decrease in ionisation energy between nitrogen and oxygen we will examine their electron configurations: Nitrogen has the electronic configuration 1s2 2s2 2p3 Oxygen has the electronic configuration 1s2 2s2 2p4

In oxygen, the electron is removed from a doubly occupied p-orbital. An electron in a doubly occupied orbital is repelled by the other electron and requires less energy to remove than an electron in a half -filled orbital. The same explanation applies for the decrease in ionisation energy from phosphorus to sulfur, except that the electrons are being removed from the 3p sub-level.

electron affinities in terms of metals and nonmetals

Metals tend to have low electron affinities and non-metal elements have higher electron affinities. This is because we are progressing towards the filling of the outer valence shell of the atom. A group 17 atom releases more energy than a group 1 atom on gaining an electron; on gaining an electron it achieves a filled valence shell and is more stable.

oxides of period 3

Na, Mg= metal = basic Al = amphoteric Si,P,S,Cl = acidic

s,p,d,f blocks

S (2 electrons)=first two groups including hydrogen and helium. P(6 electrons)=last 6 groups. D (10 electrons)=3-12 (transition metals) F (14 electrons)=rare earth metals bottom 2 groups below table,

trends in melting point for period 3

Sodium, magnesium and aluminium are all metallic elements. The strength of the metallic bond increases from sodium to aluminium, therefore, we see an increase in the melting point at the start of the period. The melting point reaches a peak with silicon, in the centre of period 3, which has the highest melting point because of its giant covalent structure. The elements to the right of silicon are all non-metallic elements. Although the covalent bonds within each molecule are strong, there are only relatively weak intermolecular forces between the molecules. Therefore, molecular covalent substances have lower melting points than giant covalent substances.

The position of hydrogen and helium

Strictly speaking, hydrogen and helium are both s-block elements as they have their valence electrons in the s sub-level. Hydrogen is often shown as an element on its own (not associated with any group) on many versions of the table, although on the periodic table in the Chemistry data booklet, it is placed above group 1. Helium is in group 18 as it has similar chemical properties to the other noble gases.

Exam tip

Students are required to be able to write equations for the reactions of Na2O, MgO, P4O10 and the oxides of nitrogen and sulfur with water. These equations are shown above as well as in section 8.5.1.

Basic oxides

The acid-base properties of these oxides are demonstrated by their reactions with water. Basic oxides, such as those shown below, dissolve in water to produce basic solutions. Na2O (s) + H2O (l) → 2NaOH (aq) MgO (s) + H2O (l) → Mg(OH)2 (aq)

The general trend down a group is decreasing electron affinity.

The additional electron gained is entering an energy level further from the nucleus. This added electron has a weaker attraction to the nucleus and therefore releases less energy when added.

The electron configuration of an element can be deduced from its position on the periodic table and vice versa.

The element lithium, for example, is in group 1 and period 2 of the periodic table. From this, we can deduce that an atom of lithium has one electron in the 2s sub-level and two occupied main energy levels. It is located in the s-block because the electrons are being added to the s sub-level. Its electron configuration is 1s2 2s1 or [He] 2s1. Iodine is in period 5 and group 17; it has seven electrons in its valence shell and five occupied main energy levels. Its electron configuration is: 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s2 5p5

Important

The group number for elements in groups 1 and 2 and 13 to 18 can be used to deduce the number of electrons in the valence shell. For example, all group 1 elements have one valence electron and all group 15 elements have 5 valence electrons. For groups 13 to 18, subtract 10 from the group number to find the number of valence electrons in the outer shell.

halogen oxidizing agents

The halogens are strong oxidising agents, with fluorine being the strongest oxidising agent and iodine the weakest (Figure 3). Simply put, the more reactive halogens are stronger oxidising agents and down the group, they become weaker oxidising agents.

First ionisation energies generally increase across a period and decrease down a group in the periodic table.

The increase in ionisation energy across a period is due to the increase in nuclear charge and decrease in atomic radii. This results in an increased attraction between the nucleus and the valence electrons. The decrease in ionisation energy down a group is due to the increase in atomic radius. The increased distance results in a weaker attraction between the nucleus and the valence electrons of an atom therefore it would require less ionization energy to remove the electron.

metallic character

The metallic character of an element is related to its ionisation energy, which depends on the nuclear charge and atomic radius of an atom. The higher the nuclear charge and smaller the atomic radius of an atom, the lower its metallic character, and vice versa. Metals tend to lose electrons in chemical reactions, as indicated by their low ionization energies.

metal oxides

The period 3 oxides show a gradual trend of decreasing metallic character across the period. Metal oxides such as sodium oxide (Na2O) and magnesium oxide (MgO) have giant ionic structures. These are solids under standard conditions because of the strong electrostatic attractions between the ions, a feature of ionic bonding.

Group

Vertical column in the periodic table

trends of atomic radii across a period

When moving across a period, the number of protons (the nuclear charge) increases by one with each successive element. The number of electrons also increases by one with each successive element. The extra electron occupies the same main energy level, which means that the electron shielding remains more or less constant across a period. This results in a stronger attraction between the nucleus and the outer (valence) electrons which pulls them closer to the nucleus. Therefore, the atomic radius decreases across the period.

Predict if a displacement reaction will occur in the following examples: a) Iodine water, I2 (aq), is added to a solution of potassium chloride, KCl (aq). b) Bromine water, Br2 (aq), is added to a solution of sodium iodide, NaI (aq).

a) No reaction occurs because iodine is a weaker oxidising agent than chlorine and cannot displace the chloride ions from solution. b) A displacement reaction occurs because bromine is a stronger oxidising agent than iodine and can displace the iodide ions from solution: Br2 (aq) + 2NaI (aq) → 2NaBr (aq) + I2 (aq) Ionic equation: Br2 (aq) + 2I− (aq) → 2Br− (aq) + I2 (aq)

endothermic processes

absorb heat

for an atom to be ionized

an electron has to be completely removed from the atom therefore the atom is bombarded with electrons therefore is endothermic must be in gaseous state therefore isolated from other attractions to allow all the energy to be used in removing an electron.

Lanthanides and Actinides

are metallic elements that make up the f-block. These two rows of elements are separated from the main body of the table, in what is known as the 'short form' of the periodic table. This form of the table is merely a matter of convenience; the table could be drawn in its 'long form'

The covalent radius is measured

as half the distance between two neighboring nuclei. covalent meaning they share electrons therefore they are merged. The other two types of measurement are the metallic radius and the van der Waals' radius.

Note that the ionic radius increases

as the number of protons decreases.

overall trends of atomic radii

atomic radii decrease across period atomic radii increase down a group

monatomic

consisting of one atom

The group number of an element can be used to.....

deduce the number of valence electrons in an atom of that element. group 13 elements have three electrons in their valence shells; group 14, four; group 15, five; group 16, six; group 17, seven and group 18, eight electrons in their valence shells.

melting point

depends on its structure and bonding. Across a period, the structure and bonding gradually change from metallic to giant covalent to molecular covalent.

Acidic oxides

dissolve in water to produce acidic solutions according to the equations below. Silicon dioxide (SiO2) does not react with water because of its giant covalent structure, but does react with strong bases such as sodium hydroxide (NaOH). P4O10 (s) + 6H2O (l) → 4H3PO4 (aq) SO3 (g) + H2O (l) → H2SO4 (aq) 3NO2 (g) + H2O (l) → 2HNO3 (aq) + NO (g) Cl2O7 (l) + H2O (l) → 2HClO4 (aq) Cl2O (g) + H2O (l) → 2HClO (aq)

intermolecular forces

forces of attraction between molecules

Non-metallic elements tend to

gain electrons to form negatively charged ions (anions), examples such as S2− and Cl−. Ions such as these are described as being isoelectronic.

Exothermic processes

give out heat

Halogens

group 17 Contains nonmetals, 7 valence electrons in it's outermost energy level. Very reactive The most reactive group of non-metals.

metallic radius

half the distance between the nuclei of two adjacent, identical metal atoms

Isoelectronic species

have the same electron configuration. Examples include cations such as Na+, Mg2+ and Al3+ and anions such as N3−, O2− and F−. These ions have different numbers of protons therefore different radii, as they are different elements, but the same number of electrons.

Metals are characterised by

having fewer valence electrons, larger atomic radii, lower electronegativity values and lower ionisation energies. They have a tendency to lose their valence electrons relatively easily to form positive ions (cations).

Non-metallic elements are characterised by

having more valence electrons, smaller atomic radii, higher electronegativity values and higher ionisation energies. They tend to gain electrons to form negative ions (anions) instead of losing electrons to form positive ions.

period

horizontal row in the periodic table

Where are metaloids located on the periodic table?

in a diagonal 'staircase' that forms a boundary between the metals and the non-metals

displacement reaction halogens

in which the more reactive halogen displaces the ions of the less reactive halogen from solution. aka switch Displacement reactions of the halogens are accompanied by distinctive colour changes. When chlorine is added to a colourless solution of potassium bromide, the colour of the solution changes to brown because of the formation of aqueous bromine. ions are colorless therefore the least reactive color shows

First electron affinity values generally

increase, becoming more exothermic, across a period.

Electronegativity

is defined as the attraction of an atom for a bonding pair of electrons.

First electron affinity

is the energy released when one mole of electrons are added to one mole of gaseous atoms to form one mole of gaseous 1− ions. First electron affinities are exothermic, whereas second electron affinities tend to be endothermic, as in the case of oxygen.

1st ionization energy

is the energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous +1 ion.

nuclear charge

is the total charge of all the protons in the nucleus of an atom. It is the same as the atomic number of the element.

van der Waals' radius

larger than covalent radii atoms dont converge measured by the distance between the nucleus

Metallic elements tend to

lose electrons from their atoms to form positively charged ions (cations). Examples include Na+, Mg2+ and Al3+. Non-metallic elements,

Positively charged ions are smaller than their parent atom.

lost their outer shells of electrons and now have one less occupied main energy level than their parent atoms. By doing so, they have become isoelectronic. Each ion still has the same number of protons in the nucleus as its parent atom, but they are pulling on fewer electrons, which increases the attraction between the nucleus and the valence electrons. For these reasons, positive ions are smaller than their parent atoms. The trend across a period is that the ionic radii decrease as the nuclear charge increases. Hence, going across period 3 from Na+ to Si4+, the ionic radius decreases.

Amphoteric

means that a substance can act as both an acid and a base.

alkaline earth metals

metallic elements in group 2 of the periodic table which are harder than the alkali metals and are also less reactive Reactive metals - but less reactive than group 1.

The majority of elements in the periodic table are

metals

Anions

negatively charged ions gain electrons

When you are asked to explain the difference between the first ionisation energies of elements in a period, remember to mention:

nuclear charge, atomic radius, and electron shielding.

Elements in the periodic table are arranged in order of increasing .....

nuclear charge, that is, the number of protons in the nucleus of an atom. This is also known as the atomic number (Z), or proton number of the element

Electron shielding

occurs when outer electrons are shielded from the attraction of the nucleus by inner electrons (known as shielding electrons) therefore the valence electrons feel less attraction because of the shielding electrons.

Group 1 elements all have

one electron in their valence shells.

Cations

positively charged ions lose electrons cats are positive and small (smaller radii than parent atom)

ionic radii

radius of an ion (an atom that has lost or gained electron(s) giving it a positive or negative charge)

periodicity

repeating patterns of chemical and physical properties.

Metalliods

share properties of both metals and non-metals

Covalent

sharing of electrons

acid base properties of oxides trend

show a change from basic to acidic across the period. Metal oxides such as magnesium oxide and sodium oxide form magnesium hydroxide and sodium hydroxide respectively when reacted with water. Aluminium oxide is amphoteric, it can act as both an acid and a base, although it does not react with water. The remaining non-metal oxides (with the exception of silicon dioxide) form acidic solutions when reacted with water.

atomic radii

size of an atom

Non-metal oxides

such as those formed by phosphorus, sulfur and chlorine are molecular compounds and exist as individual molecules. Due to the weak intermolecular forces between their molecules, they are usually gases or liquids under standard conditions.

Note #1

that all electrons in the highest occupied main energy level of an atom (the valence shell) are considered to be valence electrons, even though they may be in different sub-levels. For example, fluorine has the electron configuration 1s2 2s2 2p5. The valence electrons in the fluorine atom are those located in the 2s and 2p sub-levels. Therefore, fluorine is said to have a total of 7 valence electrons; 2 electrons in the 2s sub-level and 5 electrons in the 2p sub-level.

The first ionisation energy is

the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions. It can be thought of as a measure of the attraction between the positively charged nucleus and the negatively charged outer valence electrons. the amount of energy it takes to remove the electron....

The terms exothermic and endothermic refer to

the enthalpy change of a process.

The elements in group 17 have

the highest electron affinities (most exothermic).

In the periodic table, electronegativity increases across a period due to

the increase in nuclear charge (the number of protons in the nucleus of an atom) and the decrease in atomic radius (Figure 1). This results in a stronger attraction between the nucleus and the bonding electrons.

halide ions

the ions that are produced when atoms of chlorine and other halogens gain electrons

trends of metallic character

the metallic character decreases across a period (from left to right). Down a group, the metallic character of the elements increases.

the period number also correlates to...

the number of energy level. (n= ) the number of occupied main energy levels in the atom. Atoms of these elements have electrons that occupy the first main energy level (n = 1). Atoms of elements in period 2 such as sodium, boron and fluorine have two occupied main energy levels, n = 1 and n = 2. Atoms of elements in period 3 have three occupied energy levels and so on.

effective nuclear charge

the positive charge that an valence electron experiences from the nucleus, equal to the nuclear charge but reduced by any shielding or screening from any intervening electron distribution

oxide bonding trends

the type of bonding in the period 3 oxides changes from ionic to covalent across the period and is determined by the difference in electronegativity between the bonding atoms. Na2O, with the largest difference in electronegativity, forms an ionic compound, whereas Cl2O is covalently bonded because of the small difference in electronegativity. note they come closer to the value for oxygen (3.4) as we move to the right.

The difference in electronegativity between the atoms in a compound determines....

the type of bonding that occurs. As a general rule, a difference of up to 1.7 units results in the formation of a covalent bond, whereas a difference of 1.8 units or greater results in the formation of an ionic bond.

bromine water was added

to chlorine = clear ; no d to bromine = clear; no d to iodine = orange; d

chlorine water added

to chlorine = clear; no d to bromine = yellow; d to iodine = orange; d

iodine water was added

to chlorine = clear; no d to bromine = clear; no d to iodine = clear; no d

Group 2 elements have

two electrons in the valence shells

trends within alkali metals

very reactive group of metals The alkali metals are soft, so soft in fact that they can be easily cut with a knife. The first three members of the group, lithium, sodium and potassium, have low enough densities that they float on water. The melting and boiling points of the alkali metals are also relatively low and decrease down the group. This is due to the metallic bond getting weaker as the ionic radii of the metals increase. Note that caesium has the lowest melting point of the group (with the exception of francium) because of its large ionic radius and weak metallic bonding. Being in group 1, the alkali metals have one electron in their valence shell, which they lose relatively easily (because of their low ionisation energies) to form 1+ ions. They undergo vigorous reactions with water to form a metal hydroxide and hydrogen gas. The resulting solution has a high pH value, hence the name alkali metals. The reactivity of the alkali metals increases down the group, with lithium being the least reactive and caesium the most reactive. Francium, at the bottom of group 1, is the second rarest element on the periodic table

trends with halogens

very reactive group of non-metal elements. Note that the physical state changes down the group from chlorine, which is a gas under standard conditions, to bromine, which is a liquid. The change of state is related to the increasing molar mass of the halogens, which results in stronger intermolecular forces between the molecules. Another important point to note is that the halogens are diatomic, which means they exist as two atoms bonded together, rather than separate atoms. The melting and boiling points increase down the group, which is related to the increasing strength of the intermolecular forces the reactivity of the halogens decreases down the group. gas to solid