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Chapter 6: Periodic Trends. Development of the Periodic Table. Dmitri Mendeleev designed the first periodic table in 1869 by grouping elements with similar chemical & physical properties in rows according to Atomic mass Henry Moseley rearranged the table in 1913 according to atomic number.
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Chapter 6: Periodic Trends
Development of the Periodic Table • Dmitri Mendeleev designed the first periodic table in 1869 by grouping elements with similar chemical & physical properties in rows according to Atomic mass • Henry Moseleyrearranged the table in 1913 according to atomic number
The Modern Periodic Table Organized by groups (columns) and periods (rows)
Classifying Elements by Electron Configurations • Representative Elements- Outermost s and p are filling • Alkali metals – Outermost s has 1 e- • Alkaline earth metals – Outermost s is full • Halogens – Outermost s is full, outermost p has 5 e- • Noble Gases- Outermost s and p are totally full • Transition Metals -Outermost s is full and d is filling or full • Inner Transition Metals- Outermost s is full and f is filling or full
Periodic Trends As you move across the periodic table, you keep periodically running across the same properties... • The location of an element on the Periodic Table gives a lot of information about key properties of that element compared to the other elements Hence the name periodic table.
Trend #1: Atomic Size • An atoms size is determined by its atomic radius • Atomic radius is defined as half the distance between the nuclei of two like atoms r
Atomic Size Trend • Going down a group, atoms get bigger because the number of energy levels increases. • Going across a period, atoms get smaller. No new energy levels are added, but more protons and electrons are. The increase in protons and electrons result in a greater pull towards the nucleus. (much like magnets)
Trend #2: Atomic Size of IONS • Cations (positive ions) are always smaller than their parent atom because they lose electrons and an energy level. A strong attraction forms between the electrons left over and the nucleus. • Anions (negative ions) are always larger than their parent atom. Gaining electrons causes less attraction to the nucleus – the new electrons aren’t as attracted and are free to move around!
Answer this . . . Which is larger: S or S-2. Why? • S-2 is larger – Gaining electrons causes less of an attraction to the nucleus – the new electrons are free to move around. Which is smaller: Fe or Fe+4. Why? • Fe+4 is smaller – Losing electrons causes more of an attraction to the nucleus – the remaining electrons are pulled in tighter. Which is smaller: Na+ or Al3+? Why • They both lose an energy level BUT Al3+loses MORE e-, nucleus pulls in the remaining electrons due to the more drastic change.. Which is smaller: Be2+ or Na+. Why? • They both lose an energy level BUT Be2+ loses MORE e- & goes down to a smaller energy level. A neutron walks into a ‘restaurant’ and says, "Hey bartender give me a drink."The bartender gives him one and says, “No charge for you" An atom walks into a ‘restaurant’and proclaims, "Hey! Somebody just stole one of my electrons!"The bartender says, "Are you sure?“ The atom replies, "Yes - I'm positive!"
Trend #3: Electronegativity ELECTRONEGATIVITY - the ability for an element to attract an electron Going down a group, electronegativity decreases because the added energy levels ‘shield’ the power of the nucleus to attract electrons Going across a period, electronegativity increases because nuclear charge increases - the closer an atom is to having a full outer shell, the greater the desire to completely fill that shell by gaining electrons.
Which element would be the most electronegative? Why? • Fluorine - It is located at a low energy level with a high number of electrons and protons. There is a strong attraction between the electrons and the nucleus.It wants an electron BADLY!
How does ‘Shielding’ work? • SHIELDING EFFECT: The process of the inner electrons shielding (repelling) the outer electrons - it causes the outer electrons to be less attracted to the nucleus. • Shielding increases as you go down the periodic table because more electrons in more energy levels are added • Shielding remains constant as you go across because all the electrons in a period are in the same energy level.
Why aren’t noble gases assigned electronegativity numbers? • They are happy the way they are – they don’t need any more electrons
Trend #4: Ionization Energy IONIZATION ENERGY- The energy needed to REMOVE an electron from the outside shell. Going down a group, ionization energy decreases because more energy levels are added. This ‘shielding’ causes the electrons to be less attracted to the nucleus - little energy is needed to remove them. Going across a period, ionization energy increases because the electrons are closer to the nucleus. More energy is needed to remove them the closer they are to the nucleus.
First vs. Second • First ionization energy is the energy required to remove the first electron • Second ionization energy is the energy required to remove the second electron (and so forth) • First ionization energy is always smaller (compared to the second or third ionization energy) because each successive electron removed is closer to the nucleus and strongly attracted to it.
Trend #5: Metallic PropertiesFirst off, what’s meant by ‘metallic’?
Metallic properties increase as you go down a group, and decrease as you go across a period. X Fr is the most metallic!