A single bond is a free hinge
From the naming guides you already see an alkane as a chain of carbons, each one sp3 and pointing its four bonds toward the corners of a tetrahedron. Here is the surprising part the rung promised: that chain is not a rigid sculpture. The bond joining any two carbons is a single bond — a sigma bond — and a sigma bond is cylinder-shaped, the same all the way around its axis. Because nothing about the bond cares which way the ends are turned, the two carbons can spin relative to each other like the two halves of a hinge. Nothing breaks, nothing stretches; the molecule just twists.
Each frozen snapshot of the molecule along that twist — one particular set of angles between the bonds on the two ends — is called a conformation. The whole study of which conformations a molecule visits, and which it prefers, is conformational analysis. The crucial honesty here: conformations are not different molecules, and they are not even isomers. They are the same molecule wearing the same connectivity, caught mid-wiggle in a different pose. At room temperature the twisting is ceaselessly fast, so a real sample is a swarm of molecules constantly flickering through every conformation — but, as we will see, lingering longer in some than others.
Looking straight down the bond: the Newman projection
To reason about a twist you need a viewpoint that makes the twist visible, and the skeletal drawings from the last guide do not. The trick is to aim your eye straight down the C-C bond in question, so the front carbon sits right in front of the back carbon and the bond itself shrinks to a single point you are staring into. This view is the Newman projection. The front carbon is drawn as a dot with three lines radiating out to its other three bonds, like a three-spoke wheel. The back carbon, hidden directly behind it, is drawn as a large circle with its own three lines coming off the rim. The angle between a front spoke and a back spoke — the dihedral angle — is exactly the twist we care about.
- Pick the C-C bond you want to study and imagine standing on its axis, looking through the front carbon toward the back one — the bond runs straight away from your eye and disappears to a point.
- Draw the front carbon as a dot with three lines (a Y or three-spoke wheel) reaching out to its other three bonds.
- Draw the back carbon as a big circle behind the dot, with its own three lines coming off the rim — never letting a back line touch the centre, since that carbon is hidden behind the front one.
- Now read off the twist: if back spokes nestle in the gaps between front spokes (60 degrees offset) it is staggered; if a back spoke hides right behind a front spoke (0 degrees) it is eclipsed.
Read a Newman projection like a clock seen end-on. Spin the back circle while the front dot holds still, and you spin through every conformation of that one bond. The power of the picture is that strain — the discomfort that makes one twist worse than another — shows up purely as how close the front spokes come to the back spokes. When front and back lines lie right on top of each other, atoms are crowding; when they sit neatly between each other, everyone has room. That single visual is the whole reason chemists bother to draw it.
Ethane: staggered beats eclipsed
Ethane, CH3CH3, is the cleanest case: just two carbons, three plain hydrogens on each. Spin the back carbon and two special arrangements stand out. In the staggered conformation the three back hydrogens nestle exactly in the gaps between the three front hydrogens — a 60 degree dihedral angle, every spoke alone. In the eclipsed conformation each back hydrogen hides directly behind a front hydrogen — a 0 degree angle, the spokes paired up. These two, plus everything between, are the staggered and eclipsed conformations, and they are not equally happy. Staggered ethane sits about 12 kJ/mol lower in energy than eclipsed, which means at any instant the molecules strongly favour staggered.
Why is eclipsed worse? The hydrogens on ethane are tiny and never actually bump into each other, so it is not mainly about them taking up space. The real culprit is torsional strain: when the bonds line up eclipsed, the pairs of bonding electrons in the front C-H bonds and the back C-H bonds are forced close and repel one another. Spread the bonds out into the staggered gaps and that electron-electron repulsion eases. So torsional strain is the cost of lining bonds up, and it is highest at the eclipsed pose and lowest at the staggered pose. It is the dominant effect for a molecule as small as ethane.
Butane: when atoms actually get in each other's way
Step up to butane, CH3CH2CH2CH3, and look down the middle bond, the one between carbons 2 and 3. Now each end carries not just hydrogens but a CH3 group — a bulky chunk that does take up real space. Spinning the back carbon now gives conformations that differ in two ways at once: how eclipsed the bonds are, and how close the two methyl groups come. Two staggered forms get their own names. The anti conformation puts the two methyls on opposite sides, a 180 degree dihedral angle between them — as far apart as they can get. The gauche conformation is also staggered but sits the two methyls only 60 degrees apart, side by side.
Both anti and gauche are staggered, so both are spared torsional strain. Yet they are not equal: the gauche conformation sits about 3.8 kJ/mol higher than anti. The new cost is steric strain — two atoms or groups crowded so close their electron clouds repel, the plain physical squeeze of bulk against bulk. In gauche butane the two methyl groups are jammed near each other and complain; swing one methyl all the way around to anti and the strain melts away. So the ladder of butane conformations is: anti lowest (staggered, methyls far), gauche a little higher (staggered but crowded methyls), then the eclipsed forms higher still, with the worst being the one where the two methyls eclipse each other head-on.
Butane, down the C2-C3 bond, energy from low to high: ANTI gauche eclipsed (H/CH3) ECLIPSED (CH3/CH3) methyls methyls bonds line up methyls line up 180 apart 60 apart + crowding worst of both -------- -------- ------------ ---------------- 0 (base) ~3.8 kJ/mol ~16 kJ/mol ~19 kJ/mol torsional strain rises as bonds eclipse ; steric strain rises as the two bulky methyls are forced together.
Strain, energy, and what it does and does not change
Pull the two effects apart cleanly, because students blur them constantly. Torsional strain is about bonds lining up — it is at its worst in any eclipsed pose, even when the eclipsing atoms are tiny hydrogens, and it comes from electron pairs in adjacent bonds repelling. Steric strain is about bulk colliding — it appears only when two genuinely large groups are forced close, regardless of whether the bonds are eclipsed. Ethane shows pure torsional strain (its hydrogens are too small to clash); gauche butane shows pure steric strain (it is staggered, so no torsional penalty, yet the methyls still crowd). Most real conformations mix both.
Now connect this to energy properly. A lower-strain conformation is a lower-energy conformation, and the molecule spends more of its time in lower-energy poses — that is the Boltzmann idea behind the whole rung. But conformations are separated by tiny bumps, not real walls: the anti-to-gauche barrier in butane is only a few kJ/mol, so the population is a moving average, mostly anti with a healthy minority of gauche, all interconverting billions of times a second. This is the central honest caveat: you can never isolate a bottle of pure gauche butane. The shapes blur into one another far too fast.