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periodic table Flashcards

1
Q

what is ionisation energy

A

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

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2
Q

what’s the first ionisation energy equation for na

A

na—-na+e-

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3
Q

what happens when an electrons lost from the highest energy level

A

it w experience the least attraction from the nucleus

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4
Q

what are three factors that affect ionisation energy

A

atomic radius-the greater the distance between the nucleus and the outer electrons the less the nuclear attraction.The force of attraction falls off sharply with increasing distance so atomic radius has a large effect.

nuclear charge-the more protons there are in the nucleus of an atom the greater the attraction between the nucleus and the outer electrons.

electron shielding-electrons are negatively charged and so inner-shell repulsion repel outer shell electrons.This repulsion is the attraction between nucleus and outer electrons

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5
Q

what’s mg’s second ionisation energy and what’s definition for sIE

A

mg+—mg2+ +. e-
The energy required to remove one electron from each ion in one mole of gaseous 1+ ions of An element to form one mole of 2+ ions

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6
Q

is the second ionisation energy greater than first

A

yes as after the first electron it becomes a + ion the second electron is pulled closer to the nucleus n held more tightly.The positive ion has less electron–electron repulsion.
The nucleus now has a stronger pull on the remaining electrons..The nuclear attraction on the remaining electron increasing and more ionisation energy needed to remove the second electron

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7
Q

What do successive ionisation energies provide evidence for

A

Different electron energy levels in an atom

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8
Q

On an ionisation energy chart for fluorine the spots higher up with a higher ionisation number have more ionisation energy why?

A

They are the ones closer to the nucleus and it changes from a n-2 to n-1 so the the large increase between 7 and 8 shows that the 8th electron must be removed from a diff shell closer to nucleus w less shielding

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9
Q

What do successive ionisation energies allow perfections to be made about

A

Number of electrons in outer shell
Group of element in periodic table
Identity of element

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10
Q

1-578
2-1817
3-2745
4-11577
5-14843
6-18389
Explain this ionisation chart

A

Huge inc between 3 and 4 shows 4th electrons being removed from an innner shell so there must be three electrons in the outer shell and the element must be in gotuo 3

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11
Q

What do periodic trends in first ionisation energy provide

A

Evidence for the existence of shells and subshells

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12
Q

What happens w first ionisation energy down a group

A

Decreases due to atomic radius increases so there more inner shells so shielding increases so nuclear attraction on outer electrons decreases.

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13
Q

What’s the trend in first ionisation energy across a period?

A

nuclear charge increases as no of protons inc. electrons added to same shell so there’s similar shielding, nuclear attraction increases causing atomic radius to decrease so first ionisation energy increases

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14
Q

In period 2 and 3 first ionisation energy falls in two places why?

A

Due to the existence of subshells, their energies and how programs fill with orbitals

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15
Q

Compare Beryllium and boron

A

The 2P subshell in boron has a higher energy than the 2s subshell in beryllium. So the 2P electron is easier to remove than one of the 2S electrons in beryllium. So the first ionisation energy in boron is less than first ionisation energy of beryllium

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16
Q

Compare nitrogen and oxygen

A

So there’s a fall from nitrogen to oxygen due to electron pairings in p orbital.
In nitrogen and oxygen highest energy electrons are in 2P subjsells
In oxygen the paired electrons in one of the 2P orbital’s repel each other making it easier to remove an electron from an oxygen atom than nitrogen.So ionisation energies of oxygen is less than nitrogen
In nitrogen there’s one electron in each box so equal repulsion as far as possible
In oxygen paired electronss repel

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17
Q

What’s a semi metal/metalloid

A

Elements near the metal/non metal divide that show inbetween properties

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18
Q

What metal isn’t a solid at rtp

A

Mercury

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19
Q

What’s metallic bonding and describe the structure

A

The strong electrostatic force of attraction between cations and delocalised electrons

Del electrons spread throughout whole structure as each atom has donated its vestige or electrons to shared pool of electrons. They are mobile throughout the whole structure. It surround cations which are fixed in position keeping structure and shape of metal. They are held together in a giant metallic lattice

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20
Q

Properties of metals

A

Strong metallic bonds-as there’s a great attraction between positive ions and delocalised electrons.
High electrical conductivity-delocalised electrons are mobile so they can carry charge throughout the structure due to DEL electrons moving towards positive terminal whilst negative terminal donates electrons that move into structure
High Mp +Bp-string metallic bonds holding together atoms on the giant metallic lattice and large amount of energy needed to overcome strong electrostatic attraction been cations and electrons.
inSoluble -only react they don’t dissolve

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21
Q

What’s a giant covalent structure
+ simple molecular structure

A

strong covalent bonds in a continuous network of atoms diamond, silicon dioxide, graphite

Molecules that form a simple molecular lattice structure held together by weak intermolecular forces-h20 n co2

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22
Q

What structure does carbon silicon and boron have and why

A

Giant covalent lattice -where many billions of atoms are held together by a network of strong covalent bonds to form a giant covalent lattice
Carbon and silicon are in group 4 of the periodic table and have 4 electrons in outer shells.They use these 4 electrons to form covalent bonds to other Carbon or silicon atoms. This results in a tetrahedral structure with a bond angle of 109.5.

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23
Q

diamond

A

Diamond has a giant covalent structure.
Each C atom forms 4 covalent bonds in a tetrahedral lattice.
Extremely hard, high mp and non-conductive.

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24
Q

Structure and bonding in diamon(carbon)

A

See phone

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25
Properties of substances with a giant covalent lattice structure
High Mp and Bp-as covalent bonds are strong. So high temperatures needed to provide the large quantity of energy needed to break strong covalent bonds Insoluble-covalent bonds holding together the atoms in the lattice are far too strong to be by broken by interaction with solvents They’re non conductors of electricity-the only exceptions are graphene and graphite (which are forms of carbon) Diamond and silicon-all 4 outer shell electrons are involved in covalent bonding so none to conduct electricity Carbon can be used to make structures that conduct electricity though eg G AND G!!
26
What is graphene and graphite
Giant covalent structures of carbon based on planar heading layers with bond angles of 120 by electron pair repulsion. Graphene-single layer of graphite made of hexagonallt arranged carbon atoms held together by strong covalent bonds . Thinnest and strongest material ever made Graphite- composed of parallel layers of hexahonally arranged c atoms stacked up. The layers are bonded by weak London forces.Bonding only used three of carbons 4 outer shell electrons the spare electron is delocalised so charge can be carried.
27
What does graphene have the same electrical conductivity as
Copper
28
What can graphene and graphite be used for and why is it used in manufacturing industry
Graphene-electronics energy storage and sensors Graphite-pencils and lubricants Cuz it’s heat tolerant due to strong covalent bonds
29
Dot and cross diagram of graphene and graphite
See phone
30
Why are giant covalent substances stable and generally unreactive
Due to the the strong covalent bonds
31
What are the trend in melting points across period 2 and 3
Melting point increases from G1-G4 Decrease in MP between G4 and G5 Mp are low from G5 to G0 Sharp decrease in MP is due to a change from giant to simple molecular structure Giant structures have strong forces to overcome so have higher melting points. Simple molecular strictures have weak forces( London and IM) to overcome so lower melting points Trend goes across p2 to P3 and continues across s and p block from P4 downwards
32
Forces broke on boiling for Na Mg Al Si S Cl
Metallic Metallic Metallic Strong covalent bonding London forces Intermolecular
33
Compare and explain the electrical conductivities of sodium, sodium chloride and chlorine
Sodium- metallic bonding so good conductor Sodium chloride- ionic compound so when dissolved it’s a good conductor of electricity Choline-lack of free electrons to carry charge
34
Describe how group 2 elements can be a reducing agent
They have two outer shell electrons( in a subshell)two more than the electronic configuration of a noble gas. Each metal atom loses two electrons to form a 2+ion with electronic configuration of noble gas. So its oxidised x Another species gains 2 electrons and gets reduced so group 2 elements are reducing agent as it’s reduced another species.
35
Draw an equation for calcium being a reducing agent
Ca—Ca2+ +2e [Ar]4s2—-[Ar]
36
What happens when a group 2 element reacts with oxygen (magnesium in this case)
Redox reaction making a metal oxide as 2Mg+O2—-2MgO Mg goes from 0-2+ (4+) And O2 goes from0-2-(4-)
37
What happens when g2 reacts with water
Redox reaction making Alkaline hydroxide+ Hydorhwn M(OH)2 +H Water and MH react slowly but reaction becomes more vigorous further down the group as reactivity inc down group. Eg Sr +2H20—-Sr(Oh)2+H Sr 0-+2 so its Oxidation H+1—0 so its reduction (Total dec is -1 as two hydrogens decrease by 1 forming h2 but other 2 don’t change to form the Sr(Oh)2
38
What happens when g2 metals react with dilute acids
Redox reaction forming salt and hydrogen Eg Mg+2Hcl—MgCl2+H2
39
When redox reaction are carried out why does reactivity increase down group 2
Atoms of G2 elements react to form ions. Forming of these 2+ ions from gaseous atoms requires input of TWO ionisation energies which carries the energy input mostly.The total ionisation energies to form 2+ ions decreases down the hroup as elements become more reactive and stronger reducing agents M(g)—M+ + e- M+—-M2++e- Ionisation energies decrease down group as less attraction between nuc and outer e- so higher atomic radius and higher shielding
40
What happens when group 2 oxides react with water (CaO)
Release hydroxide ions and alkaline solutions of metal hydroxides CaO +H20—-Ca2+ +2OH- G2 hydroxides are only a little soluble in water so when they comes saturated the metal and hydroxide become solid Ca2+ + 2OH— ca(OH)2(s)
41
Solubility of g2 metal hydroxides
Increased down group so they increase in OH- ions and more alkaline So Mg(OH)2 is only slightly soluble in water with a low OH- concentration and PH 10 Ba(OH)2 is much more soluble in water and has a high OH concentration with a PH 13
42
How to show trend of solubility of metal hydroxides for group 2
1)Add spatula of group 2 oxide to water in test tube 2)Shake mixture,on scale there’s unsufficient water to dissolve all of metal hydroxide formed so saturated solution of each metal hydroxide formed with some white solid at bottom of test tube. 3)Measure Ph of each solution and alkalinity will deffooo increase down group
43
What’s group 2 compounds used for
Bases-G2 oxides,hydroxides and carbonates are used to neutralise acids Agriculture-calcium hydroxides added to field as lime to increase PH of acidic soils. Neutralising acid to form neutral water Medicine-treating acid indigestion as antacids. Many use Mg and CaCo3 as main ingredients and milk of magnesia is suspension of white MgOH used for acid indigestion
44
Trends of Boiling points of halogens (F-At)
At Rtp they all exist as diatomic molecules but they can change down the group state wise. F2-pale yellow gas Cl2- pale green gas Br2- red brown liquid I2- shiny grey black solid At2- not seen Increases in bp due to more electrons so there’s stronger London forces so more energy required to break the intermolecular forces so boiling point increases
45
How does reactivity of halogens work
Halogens have 7 outer shell electrons, one short of nearest noble gas.2 e-in s and 5 in p s2p5 Redox reactions occur and each halogen atom is reduced gaining one electron to form a 1- halide. Another species loses electrons to halogens atoms so the halogen is an oxidising agent as its oxidised another species.
46
How does displacement and reactivity link when it comes to halogen
Displacement reactions of halogens with halide ions show reactivity of halogens decreases down a group. To do a displacement reaction - a solution of each halogen added to aqueous solutions of other halides. If halogen added is more reactive then halogen displace halide from solution and solution changes colour Cl2,I2,Br2 in water=pale green,brown,orange I2 and br2 can look similar so u can add a cuclohexane and shake. Non polar solvents dissolve more easily with cyclohexane. Iodine. Will become a deeper violet.
47
When cl reacts w br what happens When cl reacts with I what gappens When br2 reacts w i what gappens
Orange colour from br2 formation Violet from i2 formation Violet from i2 formation
48
Why does trend in reactivity decreases
Halogens react with gaining electrons. Down group tendency to gain an electron decreases and halogen becomes less reactive Atomic radius increases so more inner shells and less nuclear attraction to capture electron from other species so reactivity decreases. Halogens become weaker oxidising agents down the group tendency
49
Why is fluorine the most reactive halogen
Strongest oxidising agent as it gains electrons from other species more readily than the other halogens
50
What is disproportionation
Redox reaction where the same element is both oxidised and reduced. Cl and H20 and Cl and cold dilute NAOH is examples of dsiproptionation reactions
51
What happens in a disproprtionation reaction when chlorine reacts with water and what the equation
HCl made and chloric acid. Bacteria’s killed by chloric acid and chlorate ions ClO-. It acts as a weak bleach too.(indicator paper can turn red from presence of two acids) colour disappears due to bleaching effect of chloric acid. Cl2(aq) + H20—HclO(aq) +Hcl
52
What happens in a disproprtionation reaction when chlorine reacts with cold dilute sodium hydroxide and write equations
This reactions limited by solubility of chlorine in water.If water contains dissolved sodium hydroxide more chlorine dissolves and another disproprtionation reaction takes place. It makes sodium chlorate NaClO. This is used as a household bleach. Cl2(aq)+2NaOH(aq)—NaClO+Nacl(aq) +H20(l)
53
Benefits and risks of chlorine use
Benefits-ensures waters fit to drink and bacteria’s killed.Not adding chlorine to water can be a greater risk.ensures we have a safe water supply Risks-It can be toxic and a respiratory irritant. It can react with organic hydrocarbons forming chlorinated hydrocarbons which app cause cancer. Quality of drinking water could be compromised and diseases like cholera could break out.
54
What happens when halide ions react with silver ions
Make silver halides Ag+(aq) + X——-Agx(s)
55
How are halide ions reducing agents
In displacement between halogen and halide ions react Halogen gains electrons and halide lose electrons so halogens are oxidising agents and halide ions are reducing agents. Reducing ability is shown when it reacts with Sulfuric acid which is a strong oxidising agent and halide ions will be oxidised to form the halogen. Chloride ions aren’t powerful enough to reduce H2SO4 Bromide ions are too powerful and can reduce h2So4 to So2 and iodide ions are even more powerful and reduce sulfur dioxide to sulfur and then h2S.
56
What are the equations for reduction by iodide ions of h2So4 to H2S
1)H+ + H2So4 +I- ——So2 +I2 + H20 2)H+ +So2 + I- ——-S+I2+H20 3)H++S + I- ——H2S +I2
57
What’s equation of chloride plus sodium bromide And ionic equation
Cl2(aq)+2Nabr(aq)—-2Nacl(aq)+Br2(aq) Cl2(aq)+2Br-(aq)—-2Cl- + Br2
58
How to test for anions
Carbonate test Sulfide test Halide test
59
What’s the equation for the carbonate test and how do you carry it out
Na2CO3(aq) + 2HNO3—-2NaNO3+CO2+ H20 1)add dilute nitric acid to the solid /solution to be tested 2)If you see bubbles then unknown compound could be carbonate 3)To probe Co2 present bubble gas through limewater Ca(OH)2 4)CO2 will react to form a cloudy white precipitate of calcium carbonate which turns the limewater cloudy Co2+Ca(OH)2(aq)—Caco3(s)+H20
60
What’s the equation of sulfate test and how do you carry it out
Ba2+(aq) + SO42-(aq)—-BaSO4(s) Add dilute hydrochloric acid and add barium chloride to sample (use barium nitrate if your doing a halide test).makes barium sulphate
61
What’s the equation of the halide test and how do you carry it out
Ag+(aq)+X-(aq)—-AgX(s) 1)add dilute nitric acid and aqueous silver nitrate (AgNo3) to aqueous halide solution 2) silver halide precipitates are different Silver chloride-white Silver bromide-cream coloured Silver iodide-yellow 3) add aqueous ammonia to distinguish between colours that are too similar.cl will dissolve in dilute aqueous nh3,br dissolves in conc aqeous ammonia,iodine insolbule in both. silver halides formed when aqueous silver ions react with aqueous halide ions
62
Is silver halide and barium sulfate soluble
No x
63
What’s the ionic equation of chloride Bromide Iodide And solubility’s of them
Ag+(aq) + Cl-(aq)—AgCl(s) soluble in nh3 Ag+(aq) + Br-(aq)—-agbr (s) soluble in nh3 Ag+(aq) + I(aq)—-AgI(s) insoluble in nh3
64
Barium meal
Check phone
65
If your testing a unionism inorganic compound you have to do the tests in the correct order why?
In carbonate test your adding dilute acid and looking for effervescence from co2 gas. Sulfide and halide ions make bubbles with acid. So carbonate can be carried out with no conclusions In sulfate test you add ba2+ and looking for a white precipitate of baSO4(s) Barium carbonate is white and insoluble so if we do Sunday test on carbonate you get a white precipitate so you have to do c test first and then s so we know no carbonates present In halide test you add a solution containing ag+ ions and can make silver carbonate precipitate and silver sulfate precipitate which are insoluble in water. So we need to rule out the carbonate and sulfate tests as possibilities.
66
If you add to analyse a mixture of ions what do you do
1) carbonate test- If bubbles seen continue adding nitric acid until it stops . All carbonate ions removed so none left to react in new tests 2)sulfate test- an excess of Ba(No3)2 any sulfate ions present will precipitate out as barium sulfate and filter solution to remove barium sulfate (don’t use bacl2 if we doing halide test) 3)add agNo3 Any carbonate or sulfate ions og present have been removed so any precipitate formed must have halide ions. Add Nh3 to confirm which halide you have
67
Why do we not use Sulfuric acid and HCL in carbonate test for mixture of ions
It contains surface ions and HCl contains chloride ions which will show up so we need to use HNO3
68
How test for cations and what’s the equation for it
Test for NH4- NH4+(aq)+OH-(aq)—-NH3(g)+H20 1)aqueous sodium hydroxide NAOH added to solution of an ammonium ion 2)ammonia gas produced and probs can’t see bubbles as ammonia is very soluble in water 3)mixture warmed and ammonia gas released 4)you can test gas with moist PH indicator paper as its alkaline it will turn paper blue

chem (7 decks)

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