Periodicity (HL + SL) Flashcards
How are the elements arranged in the periodic table?
Elements are arranged into four blocks associated with the four sub-level: s, p, d, and f.
What is a period in the periodic table?
The period number (n) is the outer energy level that is occupied by electrons. A “period” is a row in the periodic table.
What is a group in the periodic table?
Elements in the same vertical column are called groups, numbered 1 to 18. Elements in a group contain the same number of electrons in the outer energy level.
What is the name of group 1, 17 and 18?
group 1: alkali metals
group 17: halogens
group 18: noble gases
What can be determined from the element due to it’s position in the periodic table?
The number of principal energy level and the number of valence electrons in an atom can be deduced from its position in the periodic table. This can then be used to deduce the full electron configuration.
How can you determine what type of element it is from the periodic table?
The periodic table shows the position of metals, non-metals, and metalloids. Metals are on the left and in the centre of the table and non-metals are on the right (distinguished by the thick line). Metalloids, such as boron, silicon, and germanium, have properties intermediate between those of a metal and a non-metal.
What is d-block and f-block elements known for?
The d-block elements (group 3 to 12) are known as the transition metals. The f-block elements form two distinct groups: the lanthanoids (from La to Lu) and the actinoids (from Ac to Lr).
Example of how to interpret an elements electron configuration from its position on the periodic table.
An element’s position in the periodic table is related to its valence electrons, so the electronic configuration of any element can be deduced from the table, e.g. iodine (Z=53) is a p-block element. Iodine is in group 17 so its configuration will contain ns^2np^5. If one takes H and He as being first period, then iodine is in the fifth period, so n=5. -the full configuration for iodine will therefore be: 1(s^2)2(s^2)2(p^6)3(s^2)3(p^6)4(s^2)3(d^10)4(p^6)5(s^2)4(d^10)5(p^5) or [Kr] 5(s^2)4(d^10)5(p^5)
What is periodicity?
Elements in the same group tend to have similar chemical and physical properties. There is a change in chemical and physical properties across a period. The repeating pattern of physical and chemical properties shown by the different periods is known as periodicity
Where can periodic trends be seen?
These periodic trends can be seen clearly in atomic radii, ionic radii, ionization energies, electronegativities, electron affinities, and melting points.
What is the trend around atomic radius in the periodic table?
The atomic radius is the distance from the nucleus to the outermost electron. Since the position of the outermost electron can never be known precisely, the atomic radius is usually defined as half the distance between the nuclei of two bonded atoms of the same elements.
- As you go down a group, the outermost electron is in higher energy level, which is further from the nucleus, so the radius increases
- As you go across a period, electrons are being added to the same energy level, but the number of protons in the nucleus also increases. This attracts the energy level closer to the nucleus, so the atomic radius decreases across a period
What is a cation?
Cation is a positive ion.
What is a anion?
Anion is a negative ion.
What is the trend of cation and anion size down a group?
Both cations and anions increase in size down a group as the outer level gets further from the nucleus.
What are the trends for cations?
Cations contain fewer electrons than protons, so the electrostatic attraction between the nucleus and the outermost electron is greater and the ion is smaller than the parent atom. It is also smaller because the number of electron shells has decreased by one. Across a period, the ions contain the same number of electrons (isoelectronic), but an increasing number of protons, so the ionic radius decrease.
What are the trends for anions?
Anions contain more electrons than protons so are larger than the parent atom. Across a period, the size decreases because the number of electrons remains the same but the number of protons increases.
What are melting points trends dependent on?
Melting points depend both on the structure of the element and on the type of attractive forces holding the atoms together.
What are the trends of melting points across a period? (e.g. period 3)
- At the left of the period, elements exhibit metallic bonding ( Na, Mg, and Al), which increases in strength as the number of valence electrons increases.
- Silicon, the metalloid, in the middle of the period, has a macromolecular covalent structure, with very strong bonds, which results in a very high melting point.
- Elements in group 15, 16, and 17 ( Pa, S8 and Cl2) show simple molecular structures with weak forces of attraction between the molecules.
- The noble gas Ar exists as monatomic molecules (single atoms) with extremely weak forces of attraction between the atoms.
What are the trend for melting points in down a group? (e.g. group 1 & 17)
- In group 1, the melting point decreases down the group as the atoms become larger and the strength of the metallic bond decreases.
- In group 17, the attractive forces between the diatomic molecules increase down the group, so the melting points increase.
What is the first ionization energy trend down a group?
The first ionization energy values decrease down each group as the outer electrons is further from the nucleus and therefore less energy is required to remove it, e.g. for group 1 elements, Li (1s2 2p1, 520), Na ([Ne]3s1, 496), and K ([Ar]4s1, 419).
What is the first ionization energy trend across a period?
The values increase across a period because the extra electrons fill the same energy level and the extra protons in the nucleus attract this energy level more closely, making it harder to remove an electron.
What are the exception to the trend of the first ionization across a period?
The values do not increase regularly across a period because new sub-level. This explains why value for B (1s2 2s2 2p1) is slightly lower than the value for Be (1s2 2s2) and the value for Al ([Ne]3s2 3p1) is slightly lower than Mg ([Ne]3s2). There is also a drop in value between N (1s2 2p2 2p3) and O (1s2 2p2 2p4) and between P ([Ne]3s2 3p3) and S ([Ne]3s2 3p4). This is increased repulsion, so the paired electron is easier to remove compared with when the three electrons are all unpaired, one each in the three separate p orbitals.
What is the trend of electronegativity down a group and across a period?
As the atomic radius decreases, the attraction to the shared pair of electrons increases, so the electronegativity value increases across a period and down a group. The three most important electronegative elements are fluorine, oxygen, and nitrogen.
What is the trend of electron affinity?
The electron affinity is the energy change when an electron is added to an isolated atom in the gaseous state, i.e. X(g) + e^- -> X^-(g)
Atoms “want” an extra electron, so electron affinity values are negative for the addition of the first electron. However, when oxygen form the O^2- ion, the overall process is endothermic:
O(g) + e^- –> O^- (g) enthalpy change= -141 kJ mol^-1
O^-(g) + e^- –> O^2- (g) enthalpy change= +753 kJ mol^-
overall O(g) + 2e^- –> O^2- (g) enthalpy change= +612 kJ mol^-
What are the trends of the elements in group 1?
- Li, Na, K all contain one electron outer shell
- All reactive metals + are stored under liquid paraffin to prevent them reaction with air
- react by losing their outer electron to form metal ion
- readily lose an electron, they are good reducing agents
- reactivity increase down group as outer electron is in successively higher energy levels and less energy is required to remove it
Why are group 1 metals called alkali metals?
They called alkali metals because they all react with water to form an alkaline solution of the metal hydroxide and hydrogen gas. E.g. Lithium floats and reacts quietly; sodium melts into a ball which darts around on the surface; and the heat generated from the reaction with potassium ignites the hydrogen.
What is the equations for group 1 metals reaction with water?
2Li(s) + 2H2O(l) –> 2Li^+ (aq) + 2OH^-(aq) + H2(g)
Same but replace Li with any group 1 metal
What is the equation for group metals reacting with halogens?
The alkali metals also react readily with chlorine, bromine, and iodine to form ionic salts, e.g.
2Na(s) + Cl2(g) –> 2Na^+Cl^-(s) (2NaCl)
Same but replace Na with any group 1 metal and Cl2 with any halogen
What are the trends of the elements in group 17 (halogens)?
- halogens react by gaining one more electron to form halide ions
- good oxidizing agents
- reactivity decreases down the group as outer shell is at increasingly higher energy levels + further from nucleus
- together with the fact that there are more electrons between nucleus + outer shell, decreases attraction for an extra electron
What are the reaction between halogens?
Chlorine is a stronger oxidising agent than bromine, so can remove electron from bromide ions in solution to form chloride ions and bromine. Similarly, both chlorine and bromine can oxidize iodide ions to form iodine.
Cl2 (aq) + 2Br^-(aq) –> 2Cl^- (aq) + Br2(aq)
Cl2 (aq) + 2I^-(aq) –> 2Cl^-(aq) + I2 (aq)
Br2 (aq) + 2I^- (aq) –> 2Br^- (aq) + I2 (aq)
How do you test for the presence of halide ions?
The presence of halide ions in solution can be detected by adding silver nitrate solution. The silver ions react with halide ions to form a precipitate of the silver halide. The silver halides can be distinguished by their colour. These silver halides react with light to form silver metal. This is basis of old-fashioned film photography.
Ag^+(aq) + X^- (aq) –> AgX(s) –> (light) Ag (s) + 1/2 X2
where X= Cl, Br, or I AgCl white, AgBr cream, AgI yellow
What the trend of metals?
- tend to be shiny
- good conductors of heat & electricity (Mg, Na, Al)
Why is silicon a semiconductor?
Silicon is a semiconductor and is called a metalloid as it possesses some of the properties of a metal and some of a non-metal.
What is the relationship between phosphorus, sulphur, chlorine, and argon?
They are all non-metals & thus do not conduct electricity.
How can you distinguish between metals and non-metals?
Metals can also be distinguished from non-metals by their chemical properties. Metal oxides tend to be basic, whereas non-metal oxides tend to be acidic. Metallic properties increase down a group and metals and non-metals can exist in the same group. E.g. group 14 C= non-metal, Si & Ge= metalloids, Sn & Pb= metals
What is the exception to metal oxides being basic?
Aluminium is a metal but its oxide is amphoteric, that is, it can be either basic or acidic depending on whether it is reacting with an acidic or a base.
What is the effect of sulphur trioxide (SO3), sulphur dioxide (SO2), oxides of nitrogen (NO2) and carbon dioxide (CO2)?
Sulphur trioxide, SO3, reacts with water to form sulphuric acid. sulphur dioxide, and oxides of nitrogen, such as NO2, are the main gases responsible for formation of acid rain. Carbon dioxide, CO2, is also acidic. It dissolves in rain water forming carbonic acid, but carbonic acid is too weak and acid to cause acid rain, which defined as having a pH below 5.6. However, increasing atmospheric levels of carbon dioxide are responsible for ocean acidification, which affects marine life such as coral.
SO3 (g) + H2O (l) –> H2SO4 (aq)
SO2 (g) + H2O (l) –> H2SO3 (aq)
CO2 (g) + H2O (l) –> H2CO3 (aq)
What is the oxidation state of an atom or ion in a compound?
The oxidation state of an atom or ion in a compound is the number assigned to that atom to show the number of electrons transferred in forming a bond. If the compound was composed of ions, the oxidation state would be the charge on that atom. Oxidation states can be useful tool to identify which species has been oxidized and which reduced in a redox reaction. They are also useful to help determine the correct formula & name for ionic compound.
What are the rules to assigning oxidation states?
- In an ionic compound between 2 elements, oxidation state of each element is equal to charge carried by ion
- For covalent compounds, assume that compound is ionic, with more electronegative element forming negative ion
- The algebraic sum of all oxidation states in a compound equals zero
- the algebraic sum of all oxidation states in an ion equals charge on ion
- Elements not combined with other elements have an oxidation state zero
- Oxygen, when combined, always has an oxidation state of -2 except in peroxides (e.g. H2O2) when it has an oxidation state of -1
- Hydrogen, when combined, always has an oxidation state of +1 except in certain metal hydrides (e.g. NaH) when it has an oxidation state of -1
How do you name ionic compounds with metals in group 1, 2 & 3?
Metals in group 1, 2 & 3 always form positive ions (cation), with oxidation states of +1, +2 & +3, so it is not usual to indicate this in the name. E.g. sodium oxide, magnesium sulphide, aluminium chloride and calcium fluoride.
How do you name ionic compounds with transition metals?
Transition metals (& some other metals, e.g. tin & lead) form more than one ion as they have variable oxidation states. Roman numerals are used for oxidation numbers so names of two oxides of copper are copper(I) oxide, Cu2O, & copper(II) oxide, CuO.
What are oxyanions?
Oxyanion = formed when oxygen combines with another element to form a complex ion.
How do you name oxyanions?
In past, these were given diff names to distinguish between them. E.g. sulphate, SO4^-2, sulphite, SO3^-2, nitrate, NO3^-, & nitrite, NO2^-. These names are still commonly used even though, technically, they should now all have suffix “-ate”. SO4^-2 is sulphate(VI) ion to distinguish it from sulphate(IV) ion, SO3^-2. E.g. nitrate(V) & nitrate(III), NO3^- & NO2^-; manganate(VII) for permanganate ion, MnO4^-2, & manganate(VI) for manganate ion, MnO4^-2.
What is a transition element? And which elements aren’t typical transition metals?
Transition element = element that possesses an incomplete d sub-level in one or more of its oxidation states. Scandium isn’t a typical transition metal as its common ion, Sc^+3, has no d electrons. Zinc isn’t a transition metal as it contains a full d sub-level in all its oxidation states. For, Cr & Cu it’s more energetically favourable to half fill or completely fill d sub-level, respectively, so that they contain only one 4s electron.
What are the characteristics of transition metals?
- high melting point
- variable oxidation states
- form complex ions with ligands
- coloured compounds
- display catalytic & magnetic properties
What and how are variable oxidation states formed?
3d & 4s sub-level are very similar in energy. When transition metals lose electrons they lose 4s electrons first. All transition metals can show an oxidation state of +2. Some of transition metals can form +3 or +4 ions (e.g. Fe^3+, Mn^+4) as values of ionization energies are such that up to 2 d electrons as well as 4s electrons can be lost relatively easily. Mn^+4 ion is rare &, in higher oxidation states, element is usually found not as free metal ion but either covalently bonded or as an oxyanion, e.g. MnO4^-. E.g. Cr either +3 or +6, Mn +4 or +7
What are catalysts? And which elements act as good catalysts? And why is it beneficial?
Many transition elements & their compounds are very efficient catalysts: increase rate of chemical reactions. This helps to make industrial processes more efficient & economical.
Why do transition metals exhibit magnetism?
Transition metals & their complexes that contain unpaired electrons can exhibit magnetism.
What do d-block ions act as? And why?
Due their small size, d-block ions act as Lewis acid and attract species that are rich in elements.
What are ligands?
Ligands are neutral molecules or anions that contain a non-bonding pair of electrons and which act as Lewis bases.
What can ligands form?
These electron pairs can form coordination covalent bonds with metal ion to form complex ions.
What is a common ligand?
Water is a common ligand & most (but not all) transition metal ions exist as hexahydrate complex ions in aqueous solution, e.g. [Fe(H2O)6]^3+.
What can ligands be replaced by?
Ligands can be replaced by other ligands. A typical example is addition of ammonia to an aqueous solution of copper(II) sulphate to give deep blue colour of tetraaminecopper(II) ion. Similarly, if conc of HCl acid is added to a solution of Cu^+2(aq) yellow tetrachlorocopper(II) anion is formed.
What charge do all ligands have?
Each ligand has a charge of -1.
What is the coordinate number?
Number of lone pairs bonded to metal ion is known as coordinate number. Compounds with a coordination number of six are octahedral in shape; those with a coordination number of four are tetrahedral or square planar; & those with a coordinate number of two are usually linear.
What are different types of ligands?
monodentate = Water and cyanide ions are monodentate ligands as they utilize just one non-bonding pair to form a coordination covalent bond to metal ion
bidentate ligand = two non-bonding pairs
How is the colour of compounds determined?
Compounds Sc^+3, which have no d electrons, & of Cu^+1 & Zn^+2, which both have complete d sub-shells, are colourless. This strongly suggests that colour of transition metal complexes is related to an incomplete d level.
1. nature of transition element e.g. Mn^+2 (aq) & Fe^+3 (aq) both have configuration [Ar]3d5. Mn^+2 (aq) is pink whereas Fe^+3 (aq) is yellow
2. oxidation state - Fe^+2 (aq) is green whereas Fe^+3 is yellow.
3. Identity of ligand - [Cu(H2O)6]^+6 , is blue; [Cu(NH3)4(H2O)2]^+2, is blue/violet; & [CuCl]^-2 is yellow (green aqueous solution)
4. stereochemistry of complex - colour is also affected by shape of molecule or ion. E.g. [Cu(H2O)6]^+2 is octahedral whereas [CuCl4]^-2 is tetrahedral.
What are the five d orbitals called?
In free ion, 5 d orbitals are all of equal energy, so they are called degenerate.
What happens to d orbitals when ligand approaches a metal?
Ligands act as Lewis bases & donate a non-bonding pair electrons to form a coordination bond. As ligands approach metal along axes to form an octahedral complex, non-bonding pairs of electrons on ligands repel d x^2-y^2 & dz^2 orbitals, causing 5 d orbitals to split, 3 to lower energy & 2 to higher energy. Diff in energy between 2 levels, delta E, corresponds to wavelengths of visible light.
What happens when white light falls on aqueous solution complex?
When white light falls on aqueous solution of complex, colour corresponding to detla E is absorbed as an electron is promoted, & transmitted light will be complementary colour. E.g. [Cu(H2O)6]^+2 absorbs red light so compound appears blue.
What determines the colour that the complex becomes when white light is shone?
Amount that d orbitals are split determines exact colour. Changing transition metal changes number of protons in nucleus, which affects levels. Similarly, changing oxidation state affects splitting as number of electrons in levels is diff. Diff ligands also cause diff amounts of splitting, depending on their electron density.
How can ligands be arranged? And what is it called?
Ligands can be arranged in order of their ability to split d orbitals in octahedral complexes. This order is known as spectrochemical series.
What does the spectrochemical series show?
Iodide ions cause smallest splitting & carbonyl group CO, causes largest splitting. energy of light absorbed increases when, e.g. ammonia is substituted for water in Cu^+2 complexes as splitting increases, i.e. in going from [Cu(H2o)6]^+2 to [Cu(NH3)4(H2O)2]^+2. This means that wavelength of light absorbed decreases & this is observed in colour of transmitted light, which changes from blue to blue-violet (purple) colour.
What coloured does the solution appear depending on what light was absorbed?
If red/orange light is absorbed, solution appears blue-green as that is transmitted complementary colour. Complementary colours are opposite each other in this “colour wheel”.
