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Chapter 7
Families and Trends in the Periodic Table
| Question | Answer |
|---|---|
| What is the atomic radii | The atomic radius (covalent radius) of an atom is half the distance between the nuclei of two atoms of the same element that are joined together by a single covalent bond. |
| Why can't scientists just measure the distance from the nucleus to the end of the electron cloud and what is this uncertainty called | since the position of an electron cannot be precisely located (Heisenberg Uncertainty Principle), it is not possible for chemists to measure exactly where the electron cloud of an atom ends. |
| What happens to atomic radii when you go down a group | The values of the atomic radius increase down the groups in the Periodic Table |
| What are the two reasons for atomic radii increasing down groups | New energy level: As you move down a group, the additional electrons are going into a new energy level (orbit) which is further from the nucleus. With each extra energy level (orbit) the outer electrons are becoming further away from the nucleus, i.e. the atomic radius increases. Screening effect of inner electrons: Even though the nuclear charge increases going down a group, the influence of the increase in the nuclear charge is lessened because of what is called a screening effect. |
| What is screening effect | In an atom with one electron, that electron experiences the full attractive force of the positive nucleus. However, in an atom with many electrons, the electrons in the inner energy level(s) help to screen or shield the outer electrons from the positive charge in the nucleus. Therefore, even though the nuclear charge increases, the increase is counteracted by the increase in the shielding effect of the inner energy levels of electrons. |
| What is effective nuclear charge | This phrase refers to the net positive charge pulling the electrons towards the nucleus. The effective nuclear charge is obtained by subtracting the effect of the screening electrons from the effect of the positive charge of the nucleus. When the number of protons in the nucleus is increasing without any increase in the screening effect of the electrons in the inner energy levels, we say that there is an increase in the effective nuclear charge. |
| What happens to atomic radii when you go across the periodic table | The values of the atomic radius decrease across a period in the Periodic Table |
| What are the two reasons for atomic radii decreasing across a group | Increase in effective nuclear charge No increase in the screening effect |
| Explain how increase in effective nuclear charge causes the atomic radii to decrease | The number of protons in the nucleus increases from left to right across any one period in the Periodic Table. This means that there is a greater attractive force on the outer electrons. The increased nuclear charge tends to draw the energy levels closer to the nucleus. As this attractive force increases from left to right across the Periodic Table, the nucleus pulls the outer electrons closer to it. Therefore, the atoms become smaller, i.e. the atomic radius decreases. |
| Explain how no increase in screening causes the atomic radii to decrease | In moving from left to right across any one period in the Periodic Table, the extra electrons being added are going into the same outer energy level as the previous electron. Therefore, there is no increase in the screening effect caused by the electrons in the inner energy levels. Since electrons are being added to the same energy level, there is no increase in the number of inner energy levels to cause additional shielding. |
| Why might some elements be highly reactive | The reason that they are so reactive is because they lose their outer electrons very easily. |
| What is the first ionisation energy | The first ionisation energy of an atom is the minimum energy required to completely remove the most loosely bound electron from a neutral gaseous atom in its ground state. |
| What happens to the first ionisation energy when you go down a group | The values of the ionisation energy decrease down the groups in the Periodic |
| What are the two reason for first ionisation energy decreasing when going down a group | Increasing atomic radius Screening effect of inner electrons |
| How does increasing atomic radii cause the first ionisation energy to decrease | The atomic radius increases as you move down any one group. This means that the outermost electrons are becoming further away from the attractive force of the nucleus. Therefore, it becomes easier to remove an electron from the outer energy level. Therefore, the ionisation energy values decrease. |
| How does Screening effect of inner electrons cause the first ionisation energy to decrease | The screening effect exists going down groups, the outermost electrons are somewhat shielded from the attractive force of the positively charged nucleus. Therefore, the outermost electrons are easier to remove meaning the ionisation energy values decrease. |
| What happens to the first ionisation energy when you go across the periodic table | The values of ionisation energy increases across a period in the Periodic Table |
| What are the two reasons for the first ionisation energy increasing across the periodic table | Increase in effective nuclear charge Decreasing atomic radius |
| How does increase in effective nuclear charge cause the first ionisation energy to increase | Moving across a period, the number of protons in the nucleus increases. Therefore, the attraction between the nucleus and the outer electrons (effective nuclear charge) is steadily increasing meaning the electrons are held more firmly. It requires more energy to remove one of the electrons from the outermost energy level. Therefore, the ionisation energy values increase. |
| How does Decreasing atomic radius cause the first ionisation energy to increase | The atomic radius decreases from left to right. Therefore, an electron in the outermost level is becoming closer to the nucleus making it is more difficult to remove the electron due to the increased attraction between it and the nucleus. Therefore, the ionisation energy values increase. |
| What are some Exceptions to the general trend across a period and why are they exceptions | Completely filled sublevel Be = 1s2, 2s2 Mg = 1s2, 2s2, 2p6, 3s2 Completely filled s sublevels give a stable electron configuration. Therefore, removing an electron Therefore, the ionisation energy is higher. Half-filled sublevel N = 1s2, 2s2, 2p3 P = 1s2, 2s2, 2p6, 3s2, 3p3 Removing an electron from a sublevel that is exactly half filled requires more energy. Therefore, the ionisation energy is higher. |
| What happens to ionisation energy after an electron is removed | With each electron removed, the ionisation energy increases. |
| In a graph of successive ionisation energies in an element, what does it mean when there are peaks | Each peak corresponds to a change in energy level. |
| What is second ionisation energy | Second ionisation energy is the energy required to remove an electron from an ion with one positive charge in the gaseous state |
| What is electronegativity | Electronegativity is the ability of an atom to attract the shared pair of electrons in a covalent bond. |
| What happens to the value of electronegativity when going down a group | The values of the electronegativity energy decrease down the groups in the Periodic Table |
| What are the two reasons that electronegativity value decreases when going down a group | Increasing atomic radius Screening effect of inner electrons |
| How does increasing atomic radius cause electronegativity to decrease | The atomic radius increases as you move down any one group. This means that the outermost electrons are becoming further away from the attractive force of the nucleus (effective nuclear energy). Therefore, it becomes easier to remove an electron from the outer energy level. Therefore, the electronegativity energy values decrease. |
| How does the screening effect of inner electrons cause electronegativity value too decrease | The screening effect exists going down groups the outermost electrons are somewhat shielded from the attractive force of the positively charged nucleus. Therefore, the outermost electrons are easier to remove meaning the electronegativity values decrease. |
| What happens to the value of electronegativity when going across the periodic table | The values of electronegativity increase across a period in the Periodic Table |
| What are the two reasons that electronegativity value increases when going across the periodic table | Increase in effective nuclear charge Decreasing atomic radius |
| How does increase in effective nuclear charge cause electronegativity values to increase | Moving across a period, the number of protons in the nucleus increases. Therefore, the attraction between the nucleus and the outer electrons is steadily increasing meaning the electrons are held more firmly. It requires more energy to remove one of the electrons from the outermost energy level. Therefore, the electronegativity values increase. |
| How does decreasing atomic radius cause electronegativity values to increase | The atomic radius decreases from left to right. Therefore, an electron in the outermost level is becoming closer to the nucleus making it is more difficult to remove the electron due to the increased attraction between it and the nucleus. Therefore, the electronegativity values increase. |
| How reactive are alkali metals | Very reactive |
| Why are alkali metals very reactive | All of the alkali metals are very reactive elements because they all have low first ionisation energy values. They are so reactive they must be stored under oil. |
| Are alkaline earth metals reactive | All alkaline earth metals are reactive elements |
| Why are alkaline earth metals reactive | All alkaline earth metals are reactive elements because they have low first ionisation energy values, i.e. they readily lose the two electrons in their outermost energy level. |
| Which is more reactive between alkali metals and alkaline earth metals | Alkaline earth metals are less reactive than the alkali metals but still too reactive to occur freely in nature as pure metals. |
| Why does the reactivity of the alkali metals and alkaline earth metals increase down the group? | The reactivity of the alkali metals and alkaline earth metals increases down the group because: The atomic radius increases. The screening effect of inner energy levels of electrons increases. The above two factors cause the first ionisation energy to decrease down the group, therefore, the electrons in the outermost energy levels are more easily lost. Therefore, the reactivity increases down the group. |
| What group of elements has high electronegativity values | The halogens contain some highly electronegative elements. |
| What is the most electronegative element and what does it electronegativity tell us about its reactivity | Fluorine is the most electronegative element in the Periodic Table and therefore is the most reactive. It is extremely reactive because it has 7 electrons so it is really good at pulling electrons towards it due to its high electronegativity which is attracting electrons |
| If halogens are highly electronegative what does it mean for their reactivity | The halogens are very reactive. |
| Do the halogens exist in nature | The halogens do not exist free in nature, e.g. chlorine gas has to be extracted from one of its compounds such as sodium chloride. |
| Do the halogens conduct electricity | Since the halogens are non-metals, they are poor conductors of electricity. |
| Why does the reactivity of the halogens increase up the group? | The atomic radius decreases as you go up the group. There is a decrease in the screening effect, as there are fewer inner energy levels as you go up the group. The above two factors cause the first ionisation energy to increase up the group. Therefore, the electrons in the outermost energy levels are more tightly held as you go up the group. For example, fluorine is the most electronegative element in the Periodic Table. Therefore the reactivity increases up the group (attract electronsbttr) |
| Are noble gases reactive | Noble gases are all unreactive |
| Why are noble gases unreactive | Noble gases are all unreactive due to each element containing a full outer octet. |
| How good are noble gases at conducting electricity | As they are non metals, they are poor conductors of electricity. |
Created by:
21JulianM
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