Atomic Radii & Ionisation Energy Trends

B = Helium P = Sulphur H = Lower first ionisation energy G = Higher first ionisation energy

Electronegativity = Relative attraction an atom has for shared electrons Ionisation Energy = Energy required to remove an electron from an atom Atomic Radius = Distance from the nucleus to the outermost electron Effective Nuclear Charge = Net positive charge experienced by electrons

HCl = No hydrogen bonds H2O = Has hydrogen bonds NH3 = Has hydrogen bonds CO2 = No hydrogen bonds

<p>Increasing effective nuclear charge = Increases electronegativity Decreasing atomic radius = Increases electronegativity Weaker hydrogen bonding in ammonia = Lower boiling point than water Full outer lower sub-level = Higher ionisation energy</p> Signup and view all the answers

<p>Element R = Full outer lower sub-level Element S = Lower ionisation energy than R Element H = Less stable electron configuration Element G = More stable electron configuration</p> Signup and view all the answers

<p>Beryllium = Higher ionisation energy than Boron Boron = Lower ionisation energy than Beryllium Nitrogen = Higher ionisation energy than Oxygen Oxygen = Lower ionisation energy than Nitrogen</p> Signup and view all the answers

<p>Ionisation energy = Energy required to remove an electron from an atom First ionisation energy = Energy to remove the outermost electron Second ionisation energy = Energy to remove an electron from a +1 ion Stability of a filled sublevel = Greater energy required to remove an electron</p> Signup and view all the answers

<p>1s = Filled sublevel in Helium 2p = Filled sublevel in Neon 3s = Sublevel in Sodium configuration 4s = Sublevel in Potassium configuration</p> Signup and view all the answers

<p>Potassium = 4s1 Boron = 2s2 2p1 Nitrogen = 2s2 2p3 Oxygen = 2s2 2p4</p> Signup and view all the answers

<p>K+ = Lower second ionisation energy Ne = High ionisation energy O2- = Energy needed to remove electron from less stable N3- = Higher stability due to half-filled sublevel</p> Signup and view all the answers

<p>Li = Larger atomic radius than Ne Ne = Smaller atomic radius than Li B = Atomic radius smaller than Be O = Atomic radius smaller than N</p> Signup and view all the answers

<p>n=4 = Easiest to remove electron n=3 = Requires more energy to remove electron n=2 = Very stable, requires high energy to remove n=1 = Largest increase in energy to remove electron</p> Signup and view all the answers

<p>Half-filled sublevels = Extra stability Filled outer sublevel = Requires more energy to remove Partially filled outer shell = Less stability Complete electron configuration = High ionisation energy</p> Signup and view all the answers

<p>Alkali metals = Very reactive elements with low first ionisation energy Alkaline earth metals = Less reactive than alkali metals Noble gases = Inert and do not easily form compounds Halogens = Highly reactive nonmetals</p> Signup and view all the answers

<p>Decrease in electronegativity down groups = Increasing atomic radius Increase in electronegativity across periods = Increasing effective nuclear charge Increase in reactivity of alkali metals down the group = Decreasing first ionisation energy Decrease in first ionisation energy down a group = Screening effect of inner electrons</p> Signup and view all the answers

<p>2K + ½O2 = K2O Na + H2O = NaOH + ½ H2 Li + O2 = Li2O Rb + H2O = RbOH + ½ H2</p> Signup and view all the answers

<p>Sodium = I Calcium = II Bromine = VII Neon = VIII</p> Signup and view all the answers

<p>Alkali metals = Readily form ionic compounds by losing an outer electron Noble gases = Do not occur free in nature Chalcogens = Tend to gain electrons in reactions Pnictogens = Can form both ionic and covalent compounds</p> Signup and view all the answers

<p>Ionic bond = Formed by the transfer of electrons Covalent bond = Formed by sharing of electrons Metallic bond = Involves a 'sea' of delocalized electrons Hydrogen bond = Weak attraction between molecules</p> Signup and view all the answers

<p>I - Alkali metals = Group 1 II - Alkaline earth metals = Group 2 VI - Chalcogens = Group 16 VII - Halogens = Group 17</p> Signup and view all the answers

<p>Reactivity increases down alkali metals = Due to lower ionisation energy Electronegativity increases across periods = Due to increasing nuclear charge Elements do not occur free in nature = High reactivity in elemental state Oxides are formed with O2 = Ionic compounds are produced with alkali metals</p> Signup and view all the answers

<p>Alkali Metals = Lose their shine when exposed to air Halogens = Do not exist free in nature, must be extracted from compounds</p> Signup and view all the answers

<p>Alkali Metals = More reactive with water moving down the group Halogens = Reactive due to high electron attraction</p> Signup and view all the answers

<p>First ionisation energy = Energy required to remove the most loosely-bound electron Silicon's ionisation energy vs Aluminium = Silicon has a greater nuclear charge Silicon's ionisation energy vs Carbon = Silicon has a greater atomic radius Gradual increase in ionisation energy = Indicates electrons are being removed from the same shell</p> Signup and view all the answers

<p>Sharp increase at 5th electron = First electron removed from the 2nd shell Sharp increase at 13th electron = First electron removed from the 1st shell Gradual increase in energy = Electrons being removed from the same shell Line Emission Spectrum = Experimental evidence for existence of energy levels</p> Signup and view all the answers

<p>Alkali Metals = React with oxygen in air Halogens = Exist as compounds in nature</p> Signup and view all the answers

<p>Removing the 5th electron = Sharp increase signifies a new shell Removing the 13th electron = Sharp increase signifies a new energy level Gradual increases in ionisation energy = Indicates removal from the same energy level First ionisation energy trends = Varies across the periodic table based on atomic structure</p> Signup and view all the answers

<p>Carbon = Higher first ionisation energy than silicon Silicon = Has a greater nuclear charge than aluminium Aluminium = Lower first ionisation energy than silicon Fluorine = The most electronegative element</p> Signup and view all the answers

<p>Atomic Radius Increases Down a Group = Additional electrons enter new energy levels further from the nucleus. Atomic Radius Decreases Across a Period = Increased effective nuclear charge pulls electrons closer. Ionisation Energy Decreases Down a Group = Outermost electrons are shielded by inner electrons. Ionisation Energy Increases Across a Period = Electrons are pulled more strongly by the nucleus.</p> Signup and view all the answers

<p>Atomic Radius = Half the distance between the nuclei of two atoms. First Ionisation Energy = Minimum energy required to remove the most loosely bound electron. Effective Nuclear Charge = The net positive charge experienced by valence electrons. Screening Effect = Inner electrons reduce the attractive force on outer electrons.</p> Signup and view all the answers

<p>Lithium = Higher ionisation energy than potassium. Sodium = Larger atomic radius than lithium. Beryllium = Smaller atomic radius than boron. Fluorine = Higher ionisation energy than oxygen.</p> Signup and view all the answers

<p>Increased Atomic Radius = Easier to remove outer electrons. Decreased Effective Nuclear Charge = Results in lower ionisation energy. Increased Electrons in Same Energy Level = Leads to higher atomic radius down a group. Decreased Atomic Radius = Results in higher ionisation energy across a period.</p> Signup and view all the answers

<p>Group 1 Elements = Lower ionisation energy compared to Group 2. Group 17 Elements = Higher ionisation energy due to effective nuclear charge. Noble Gases = Highest ionisation energy in their respective periods. Metals = Generally have lower ionisation energies than non-metals.</p> Signup and view all the answers

<p>Increased Atomic Size = Occurs as you move down a group. Decreased Atomic Size = Occurs as you move from left to right in a period. Ionisation Energy Drop = Seen as you descend a group. Ionisation Energy Rise = Observed as you go across a period.</p> Signup and view all the answers

<p>Valence Electrons = Electrons in the outermost shell. Core Electrons = Electrons in inner shells that shield outer electrons. Electronegativity = Tendency of an atom to attract electrons. Covalent Radius = Half the distance between nuclei of bonded atoms.</p> Signup and view all the answers

<p>Rubidium = Largest atomic radius among alkali metals. Neon = Smallest atomic radius in its period. Silicon = Intermediate atomic radius in its group. Chlorine = Smaller atomic radius than bromine.</p> Signup and view all the answers