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Fluorine has larger electronegativity than hydrogen

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Fluorine has larger electronegativity than hydrogen

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Electronegativity is defined as the ability of an atom involved in a chemical bond to polarize electron density towards itself. Electronegativity tends to increase across a Period (from left to right as we face the Table) and decrease down a Group.

Incomplete electron shells tend to shield nuclear charge very imperfectly. This is manifested by the well-known contraction in atomic radii as we go from element to element across a Period from left to right.

And thus fluorine, to the right of the second Period, Z=9, is more electronegative than hydrogen, Z=1.

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These species have more than 8 electrons around the central atom. These structures are only possible when the principal quantum number is greater than n = 3 and d orbitals are available for bonding. These are most common with large central atoms bonded to small, electronegative atoms.

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Electron Affinity and the Periodic Table

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Electron Affinity and the Periodic Table

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As we go from left to right across a period EAs tend to become more negative.

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Group of the Periodic table is the same as the number of valence shell electrons (!!!) {Almost failed a problem on this}

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The size of the atom (and its covalent radius) must increase as we increase distance.

Over a period, due to increasing effective charge the radius decreases.

Down a group, electrons are added to new shells which increases radius.

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Voiceover: Let's think a little bit about the notion of atomic size or atomic radius in this video. At first thought, you might think well this might be a fairly straight-forward thing. If I'm trying to calculate the radius of some type of circular object I'm just thinking about well what's the distance between the center of that circular object and the edge of it. So the length of this line right over here. That would be the radius. And so a lot of people when they conceptualize an atom they imagine a positive nucleus with the protons in the center right over here then they imagine the electrons on these fixed orbits around that nucleus so they might imagine some electrons in this orbit right over here, just kind of orbiting around and then there might be a few more on this orbit out here orbiting around, orbiting around out here. And you might say, "well okay, that's easy to figure out the atomic radius. I just figure out the distance between the nucleus and the outermost electron and we could call that the radius." That would work except for the fact that this is not the right way to conceptualize how electrons or how they move or how they are distributed around a nucleus. Electrons are not in orbits the way that planets are in orbit around the sun and we've talked about this in previous videos. They are in orbitals which are really just probability distributions of where the electrons can be, but they're not that well defined. So, you might have an orbital, and I'm just showing you in 2 dimensions. It would actually be in 3 dimensions, where maybe there's a high probability that the electrons where I'm drawing it in kind of this more shaded in green. But there's some probability that the electrons are in this area right over here and some probability that the electrons are in this area over here, and let's say even a lower probability that the electrons are over this, like this over here. And so you might say, well at a moment the electron's there. The outermost electron we'd say is there. You might say well that's the radius. But in the next moment, there's some probability it might be likely that it ends up here. But there's some probability that it's going to be over there. Then the radius could be there. So electrons, these orbitals, these diffuse probability distributions, they don't have a hard edge, so how can you say what the size of an atom actually is? There's several techniques for thinking about this. One technique for thinking about this is saying, okay, if you have 2 of the same atom, that are- 2 atoms of the same element that are not connected to each other, that are not bonded to each other, that are not part of the same molecule, and you were able to determine somehow the closest that you could get them to each other without them bonding. So, you would kind of see, what's the closest that they can, they can kind of get to each other? So let's say that's one of them and then this is the other one right over here. And if you could figure out that distance, that closest, that minimum distance, without some type of, you know, really, I guess, strong influence happening here, but just the minimum distance that you might see between these 2 and then you could take half of that. So that's one notion. That's actually called the Van der Waals radius. Another way is well what about if you have 2 atoms, 2 atoms of the same element that are bonded to each other? They're bonded to each other through a covalent bond. So a covalent bond, we've already- we've seen this in the past. The most famous of covalent bonds is well, a covalent bond you essentially have 2 atoms. So that's the nucleus of one. That's the nucleus of the other. And they're sharing electrons. So their electron clouds actually, their electron clouds actually overlap with each other, actually overlap with each other so the covalent bond, there the electrons in that bond could spend some of their time on this atom and some of their time on this atom right over here. And so when you have a covalent bond like this, you can then find the distance between the 2 nuclei and take half of that and call that call that the atomic radius. So these are all different ways of thinking about it. Now, with that out of the way, let's think about what the trends for atomic size or atomic radii would be in the periodic table. So the first thing to think about is what do you think will be the trend for atomic radii as we move through a period. So let's say we're in the fourth period and we were to go from potassium to krypton. What do you think is going to be the trend here? And if you want to think about the extremes, how do you think potassium is going to compare to krypton in terms of atomic radius. I encourage you to pause this video and think about that on your own. Well, when you're in the fourth period, the outermost electrons are going to be in your fourth shell. Here, you're filling out 4S1, 4S2. Then you start back filling into the 3D subshell and then you start filling again in 4P1 and so forth. You start filling out the P subshell. So as you go from potassium to krypton, you're filling out that outermost fourth shell. Now what's going on there? Well, when you're at potassium, you have 19- 1, 2, 3, 4, 5, 6, 7, 8, 9, 10, 11, 12, 13, 14, 15, 16, 17, 18, 19. You have 19 protons and you have 19 electrons. Well I'll just draw those. 1, 2, 3, 4, 5, 6, 7, 8, 9, 10, 11, 12, 13, 14, 15, 16, 17, 18, but you only have 1 electron in that outermost, in that fourth shell, so let's just say that's that electron at a moment, just for visually. It doesn't necessarily have to be there but just to visualize that. So that 1 electron right over there, you have 19, yeah, you have 19 protons. So, you have some, I guess you could say Coulom force that is attracting it, that is keeping it there. But if you go to krypton, all of a sudden you have much more positive charge in the nucleus. So you have 1, 2, 3, 4, 5, 6, 7, 8- I don't have to do them all. You have 36. You have a positive charge of 36. Let me write that, you have plus 36. Here you have plus 19. And you have 36 electrons, you have 36 electrons- I don't know, I've lost track of it, but in your outermost shell, in your fourth, you're going to have the 2S and then you're going to have the 6P. So you have 8 in your outermost shell. So that'd be 1, 2, 3, 4, 5, 6, 7, 8. So one way to think about it, if you have more positive charge in the center, and you have more negative charge on that outer shell, so that's going to bring that outer shell inward. It's going to have more I guess you could imagine one way, more Coulomb attraction right over there. And because of that, that outer most shell is going to drawn in. Krypton is going to be smaller, is going to have a smaller atomic radius than potassium. So the trend, as you go to the right is that you are getting, and the general trend I would say, is that you are getting smaller as you go to the right in a period. That's the reason why the smallest atom of all, the element with the smallest atom is not hydrogen, it's helium. Helium is actually smaller than hydrogen, depending on how you, depending on what technique you use to measure it. That's because, if we take the simplest case, hydrogen, you have 1 proton in the nucleus and then you have 1 electron in that 1S shell, and in helium you have 2, 2 protons in the nucleus and I'm not drawing the neutrons and obviously there's different isotopes, different numbers of neutrons, but you have 2 electrons now in your outer most shell. So you have more, I guess you could say, you could have more Coulomb attraction here. This is plus 2 and then these 2 combined are negative 2. They're going to be drawn inward. So, that's the trend as we go to the right, as we go from the left to the right of the periodic table, we're getting smaller. Now what do you think is going to happen as we go down the period table? As we go down the periodic table in a given group? Well, as we go down a group, each new element down the group, we're adding, we're in a new period. We're adding a new shell. So you're adding more and more and more shells. Here you have just the first shell, now the second shell and each shell is getting further and further and further away. So as you go down the periodic table, you are getting, you are getting larger. You're having a larger atomic radius depending on how you are measuring it. So what's the general trend? Well if you get larger as you go down, that means you're getting smaller as you go up. You get smaller, smaller as you go up. So, what are going to be, what's going to be the smallest ones? Well, we've already said helium is the smallest. So what are going to be some of the largest? What are going to be some of the largest atoms? Well that's going to be the atoms down here at the bottom left. So, these are going to be large, these are going to be small. So, large over here, small over here and the general trend, as you go from the bottom left to the top right you are getting, you are getting smaller.

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Atomic radius trends on periodic table (video) | Khan Academy

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Voiceover: Let's think a little bit about the notion of atomic size or atomic radius in this video. At first thought, you might think well this might be a fairly straight-forward thing. If I'm trying to calculate the radius of some type of circular object I'm just thinking about well what's the distance between the center of that circular object and the edge of it.

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The first ionization energy is the energy required to form a cation with +1 charge:

X(g)-> X⁺(g) + e⁻

Values are always positive

As size increases the ionization energy decreases.

Increases over a period

Decreases down a group

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Effective Nuclear Charge and its Implications

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Effective Nuclear Charge and its Implications

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As the effective nuclear charge increases, it becomes easier to add an electron.

As electrons are removed, the remaining core electrons experience a greater effective nuclear charge and are drawn closer to the nucleus.

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Transition metals lose their s electrons before the d electrons

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Ag+, Zn2+, Cd2+

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1. Determine the Lewis structure of the molecule.

2. Determine the number of regions of electron density around an atom using VSEPR theory, in which single bonds, multiple bonds, radicals, and lone pairs each count as one region.

3. Assign the set of hybridized orbitals from Figure 5.21 that corresponds to this geometry.

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Form from the overlap of:

1. two s orbitals

2. an s orbital and a p orbital

3. two p orbitals

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orientation of orbitals and distance affects their overlap

energy of the system depends on how much the orbitals overlap

strength of a covalent bond depends on the extent of overlap of the orbitals involved

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1. an orbital on one atom overlaps an orbital on a second atom

2.  the single electrons in each orbital combine to form an electron pair

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VSEPR Theory Song - University of Pittsburgh

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The Periodic Table Song (2018 Update!) | SCIENCE SONGS

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Hydrogen forms covalent bonds

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Polar molecules tend to align when placed in an electric field with the positive end of the molecule oriented toward the negative plate and the negative end toward the positive plate (Figure 4.28). We can use an electrically charged object to attract polar molecules, but nonpolar molecules are not attracted. Also, polar solvents are better at dissolving polar substances, and nonpolar solvents are better at dissolving nonpolar substances.

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To summarize, to be polar, a molecule must:

1. Contain at least one polar covalent bond.

2. Have a molecular structure such that the sum of the vectors of each bond dipole moment does not cancel.

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For diatomic molecules, there is only one bond, so its bond dipole moment determines the molecular polarity

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μ= Qr

where Q is the magnitude of the partial charges (determined by the electronegativity difference) and r is the distance between

the charges

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1. Write the Lewis structure of the molecule or polyatomic ion.

2. Count the number of regions of electron density (lone pairs and bonds) around the central atom. A single, double, or triple bond counts as one region of electron density.

3. Identify the electron-pair geometry based on the number of regions of electron density: linear, trigonal planar, tetrahedral, trigonal bipyramidal, or octahedral (Figure 4.19, first column).

4. Use the number of lone pairs to determine the molecular structure (Figure 4.19). If more than one

arrangement of lone pairs and chemical bonds is possible, choose the one that will minimize repulsions, remembering that lone pairs occupy more space than multiple bonds, which occupy more space than single bonds. In trigonal bipyramidal arrangements, repulsion is minimized when every lone pair is in an equatorial position. In an octahedral arrangement with two lone pairs, repulsion is minimized when the lone pairs are on opposite sides of the central atom.

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From largest to smallest:
1. lone pair
>
2. triple bond
>
3. double bond
>
4. single bond

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From greatest to least:
1. lone pair-lone pair
>
2. lone pair-bonding pair
>
3. bonding pair-bonding pair

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Noble gases do not have electronegativity, they are an exception to the periodic trend

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First Year Chemistry in the School of Chemistry at the University of Sydney

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We call the individual Lewis structures resonance forms. The

actual electronic structure of the molecule (the average of the resonance forms) is called a resonance hybrid of the

individual resonance forms.

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If two or more Lewis structures with the same arrangement of atoms can be written for a molecule or ion, the actual distribution of electrons is an average of that shown by

the various Lewis structures.

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A double bond between two atoms is shorter (and stronger) than a single bond between the same two atoms.

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1. A molecular structure in which all formal charges are zero is preferable to one in which some formal charges

are not zero.

2. If the Lewis structure must have nonzero formal charges, the arrangement with the smallest nonzero formal charges is preferable.

3. Lewis structures are preferable when adjacent formal charges are zero or of the opposite sign.

4. When we must choose among several Lewis structures with similar distributions of formal charges, the structure with the negative formal charges on the more electronegative atoms is preferable.

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We can double-check formal charge calculations by determining the sum of the formal charges for the whole structure. The sum of the formal charges of all atoms in a molecule must be zero; the sum of the formal charges in an ion should equal the charge of the ion.

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The formal charge of an atom in a molecule is the hypothetical charge the atom would have if we could redistribute the electrons in the bonds evenly between the atoms..

Thus, we calculate formal charge as follows:

formal charge = # of valence shell electrons - # of lone pairs - 1/2 # of bonding electrons

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Elements in the third and higher

periods (n ≥ 3) have more than four valence orbitals and can share more than four pairs of electrons with other atoms because they have empty d orbitals in the same shell. Molecules formed from these elements are sometimes called

hypervalent molecules.

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Many covalent molecules have central atoms that do not have eight electrons in their Lewis structures. These molecules fall into three categories:

• Odd-electron molecules have an odd number of valence electrons, and therefore have an unpaired electron.

• Electron-deficient molecules have a central atom that has fewer electrons than needed for a noble gas configuration. Generally,

these are molecules with central atoms from groups 2 and 12, outer atoms that are hydrogen, or other atoms that do not

form multiple bonds

• Hypervalent molecules have a central atom that has more electrons than needed for a noble gas configuration.

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Hydrogen is almost never a central atom in the Lewis Structure (an anomaly).

As the most electronegative element, fluorine also cannot be a central atom.

Boron tends to form compounds with fewer than eight electrons

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Lewis Structure for Complex Molecules and Ions

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Lewis Structure for Complex Molecules and Ions

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For more complicated molecules and molecular ions, it is helpful to follow the step-by-step procedure outlined here:

1. Determine the total number of valence (outer shell) electrons. For cations, subtract one electron for each positive charge. For anions, add one electron for each negative charge.

2. Draw a skeleton structure of the molecule or ion, arranging the atoms around a central atom. (Generally, the least electronegative element should be placed in the center.) Connect each atom to the central atom with a single bond (one electron pair).

3. Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen), completing an octet

around each atom.

4. Place all remaining electrons on the central atom.

5. Rearrange the electrons of the outer atoms to make multiple bonds with the central atom in order to obtain

octets wherever possible.

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Because hydrogen only needs

two electrons to fill its valence shell, it is an exception to the octet rule. The transition elements and inner transition elements also do not follow the octet rule.

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Metallic properties depend on having easily removable electron. So these increase down a group and decreases over the period.

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Electron Affinity and Groups of the Periodic Table

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Electron Affinity and Groups of the Periodic Table

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Second element most often has the greatest EA.

Why?
Because the n=2 shell is relatively small resulting in large electron-electron repulsions.

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Negative values of EA means energy is released. Positive values of EA means energy is required for the electron to become negatively charged.

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An isoelectronic series is a group of atoms/ions that contain the same number of electrons.

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Binary salts, which contain only two elements, are named for the metal and non-metal which combine to form them. So, when sodium and chlorine react, they form sodium chloride. The sodium atom loses an electron to form the Na⁺ ion, and chlorine gains an electron to form the Cl⁻ ion. Anions (negatively charged) and cations (positively charged) always combine to give neutral species, i.e., the number or positive and negative charges will be equal. So, the formula for sodium chloride is NaCl, with one positive and one negative charge. The formula of magnesium chloride is MgCl ₂. Each chloride has a charge of 1- and the magnesium ion has a charge of 2+.
Write the formulas for the salts made by combining the following elements. Don't worry about subscripts, MgCl₂ would be written as MgCl2.

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Hydrogen is a special case. It can gain an electron, to form H⁻. This is called hydride and happens when hydrogen is reacting with a metal.
Hydrogen can also lose an electron to form H⁺, a proton, when reacting with non-metals.

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When individual non-metallic elements react with metals to form salts, they gain electrons, forming anions.

Group VII elements (F → I, seven valence electrons) gain 1 electron to form the 1- state.
They are given the ending -ide, like fluoride (F⁻), chloride (Cl⁻), bromide (Br⁻) and iodide (I⁻).

Group VI elements (O → Se, six valence electrons) gain 2 electrons to form the 2- state.
They are again given the ending -ide, like oxide (O²⁻).

Group V elements (N and P, five valence electrons) gain 3 electrons to form the 3- state and have the ending -ide (N³⁻ is nitride).

Carbon, with four valence electrons, can gain 4 electrons to form carbide, C⁴⁻.

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Metals lose electrons, to form cations.
Group I metals (Li → Fr, 1 valence electron) lose 1 electron.
This leaves them in the 1+ state.

Group II metals (Be → Ra, 2 valence electrons) lose 2 electrons.
This leaves them in the 2+ state.

Most of the Transition metals can exist in a number of different states.

For example, manganese, Mn, can exist in the 2+, 3+, 4+, 6+ and 7+ states.
To specify a state, it is written in Roman Numerals, directly after the metal name.
So, Mn⁴⁺ would be written as manganese(IV).

There are some Transition metals with more limited stable states.
Silver, Ag, is only stable as Ag⁺.
Zinc, Zn, is only stable as Zn²⁺.
Cadmium, Cd, is only stable as Cd²⁺.
No Roman Numerals are used in these cases, since there is only one possible state.

Aluminum, Al, is only stable as Al³⁺. The Roman Numeral is again, not needed.

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*doesn't cover everything

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Mahjong Chem -- Free Game to Practice Chemistry Knowledge!

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Mahjong Chem is a free online game to practice chemistry knowledge and skills.

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Periodic Table Game - Level One - All Elements

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Have fun learning chemisty, the elements, and the periodic table with these free online educational games.

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Most transition metal cations have 2+ or 3+ charges that result from the loss of their outermost s electron(s) first

When forming a cation, an atom of a main group element tends to lose all of its valence electrons

The charge of a cation formed by the
loss of all valence electrons is equal to the group number minus 10

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Why is Electron Affinity Usually Negative?

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Why is Electron Affinity Usually Negative?

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Valence electrons tend to have the highest energy and are the most likely to react.

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High ionization energy and high electronegativity

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result when atoms share, rather than transfer (gain or lose), electrons

often exist as gases, low-boiling liquids, and low-melting solids under normal conditions

usually formed by a combination of nonmetals

poor conductors in any state

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main-group metals: lose enough electrons to equal number of electrons in an atom of the preceding noble gas.

nonmetal ions: gain enough electrons to equal number of

electrons in an atom of the succeeding noble gas.

Moving from the far right to the left, elements often form anions with a negative charge equal to the number of groups moved left from the noble gases.

Moving from the far left to the right, main-group elements tend to form cations with a charge equal to the group number.

Predictive value decreases when moving toward the center of the periodic table

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Ionic compounds are electrically neutral because total number of positive charges of the cations equals the total number of negative

charges of the anions.

The formula of an ionic
compound must have a ratio of ions such that the numbers of positive and negative charges are equal.

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solids that typically melt at high temperatures and boil at even higher temperatures

not electrically conductive in solid form because its ions are unable to flow

can conduct electricity when molten (or dissolved) because its ions are able to move freely through the liquid

exhibit a crystalline structure and tend to be
rigid and brittle

very strong however dissolve readily in water

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When an element composed of atoms that readily lose electrons (a metal) reacts with an element composed of atoms

that readily gain electrons (a nonmetal), a transfer of electrons usually occurs, producing ions. The compound formed

by this transfer is stabilized by the electrostatic attractions (ionic bonds) between the ions of opposite charge present

in the compound.

When a metal is combined with one or more nonmetals,

the compound is usually ionic.

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What does it mean to say that the energy of the electrons in an atom is quantized?

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What does it mean to say that the energy of the electrons in an atom is quantized?

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Electrons exist on specific energy levels and cannot exist in between them. When an electron moves to another energy level, it must release or absorb very specific amounts of energy. This is what quantized means: the amount of energy absorbed or released can only be specific quantities called quanta. @

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Watch out for the positioning of the d block

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K here goes
Avogadro's Number:
Multiply a given number of atoms (photons in certain cases) by this number to get moles.

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Let's play a game or just mess around if you want

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A cation has a smaller radius than the atom it was made from.

An anion has a larger radius than the atom it was made from.

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Electronegativity generally increases going from left to right across a period and decreases going down a group for the representative elements. Use the difference in electronegativity to describe bond polarity. Don't forget that hydrogen has electronegativity equal to phosphorus and between boron and carbon.Remember, the greater the difference in electronegativity between the bonded atoms, the more polar the bond. Look at the Periodic Table.

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While all single bonds are σ bonds, multiple bonds consist of both σ and π bonds.

Between any two atoms, the first bond formed will always be a σ bond, but there can only be one σ bond in any one location. In any multiple bond, there will be one σ bond, and the remaining one or two bonds will be π bonds.

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Consider the electronegativity trends.
The classifications ionic, polar covalent, and nonpolar covalent lie on a continuum.
In general, we can say that the bond is ionic if the difference in electronegativity is
large. This occurs when the bond involves a metal and a nonmetal. At the other end of the continuum is nonpolar covalent, when the difference in electronegativity
is essentially zero. This occurs when the bond involves two identical nonmetal atoms,
as well as bonds involving hydrogen bonded to carbon, boron, or phosphorus. All the
rest of the bonds between two nonmetals are assumed to lie between the two extremes
and can be classified as polar covalent. Look at the Periodic Table

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The following ideas are important in understanding hybridization:

1. Hybrid orbitals do not exist in isolated atoms. They are formed only in covalently bonded atoms.

2. Hybrid orbitals have shapes and orientations that are very different from those of the atomic orbitals in isolated atoms.

3. A set of hybrid orbitals is generated by combining atomic orbitals. The number of hybrid orbitals in a set is equal to the number of atomic orbitals that were combined to produce the set.

4. All orbitals in a set of hybrid orbitals are equivalent in shape and energy.

5. The type of hybrid orbitals formed in a bonded atom depends on its electron-pair geometry as predicted by the

6. VSEPR theory.

7. Hybrid orbitals overlap to form σ bonds. Unhybridized orbitals overlap to form π bonds.

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group 2 (2A): have a completely filled ns subshell, electron must be added to a higher np subshell which is more difficult

group 15 (5A): have a half filled np subshell, electron must be paired with an existing np

group 18 (8A): have a completely filled shell, electron must be added to a higher n which is more difficult

Initial relative stability of the electron configuration disrupts the trend

group 17 (7A): have the most negative electron affinities